General Introduction and Nature of Matter Quick Review · Chapter - 1 : Some Basic Concepts of Chemistry TOPIC-1 General Introduction and Nature of Matter Quick Review Ø Chemistry

Chapter - 1 : Some Basic Concepts of Chemistry

TOPIC-1General Introduction and Nature of Matter

Quick Review Ø Chemistry is that branch of science which deals with the composition, structure and properties of matter. Ø Chemistry is referred to as the science of atoms and molecules. Ø Chemical plays a central role in science and is often related to other branches of science like physics, biology

etc. Ø Chemistry plays an important role in our daily life and is related to fields like food, fuels, textiles, dyes, drugs,

disinfectants, perfumes, building materials, paints, inks, fertilizers, insecticides, pesticides, soaps, detergents, alloys etc.

Ø Chemistry plays a crucial role for fulfilling needs of human beings like food, health care products and other materials aimed at improving the quality of life.

Ø Nature of matter: Matter is anything that has mass and occupies space. Ø Matter can exist in three physical states viz. solid, liquid and gas. Ø Solids have definite volume and definite shape. Ø Liquids have definite volume but not the definite shape. Ø Gases have neither definite volume nor definite shape. Ø The three states of matter are interconvertible by changing the conditions of temperature and pressure.

Solid Liquid Gas

heat

cool

heat

cool� ⇀��↽ �� ⇀��↽ ��

Ø Matter can be categorised as mixture or pure substance. Ø The matter which contains two or more substances is called mixture. Ø A mixture may be homogeneous or heterogeneous. Ø The mixture in which the composition is uniform throughout the mixture is called homogeneous mixture. e.g.

sugar solution in water. Ø The mixture in which the composition is not uniform throughout the mixture is called heterogeneous mixture.

e.g. mixture of salt and sugar. Ø Pure substances have fixed compounds. These can be classified into elements and compounds. Ø An element contains only one type of particles. These particles may be atoms or molecules. Ø A compound is a combination of two or more atoms of different elements. Ø Properties of Matter and their Measurement—The general properties of matter are categorised into two types : (i) Physical properties (ii) Chemical properties Ø Unit of Measurement—A unit of measurement is a definite magnitude of a physical quantity, defined and

adopted by convention or by law that is used as a standard for measurement of the some physical quantity. For example, length is a physical quantity. Metre is a unit of length that represents a definite pre-determined

length. Ø Some Physical Quantities— (a) Mass and Weight (i) Mass—Mass is a quantity representing the amount of matter in a particle or object. The standard

unit of mass in the International System of Units (SI) is kilogram (kg). (ii) Weight—Weight is a quantity representing the force exerted on a particle or object by an accelera-

tion due to gravitational field of the earth. In the International System of Units (SI), weight can be expressed in terms of the force, in Newton (N).

W = m × g where m = mass, g = gravitational force (b) Volume—Volume is the quantity of three dimensional space occupied by a liquid, solid and gas. SI unit

of volume is m3.

2 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI (c) Density—

l The density of a substance is the relationship between the mass of the substance and how much space it takes up i.e. volume.

l Density is equal to the mass of the substance divided by its volume.

Density = Mass/ Volume

d = m/V

SI Unit of density = Kg/m3

(d) Temperature—In everyday term, temperature is a measure of the ‘‘hotness’’ or ‘‘coldness’’ of a substance. l Temperature is the property of matter which reflects the quantity of energy of motion of the

component particle. l There are several scales used to measure this value (e.g., Kelvin (K), degree Celsius (°C), degree

Fahrenheit (°F)). (e) Length—It is the longest dimension of any object, in distinction from breadth or width, extent of anything

from end to end. l The property of being the extent of something from beginning to end can be defined as length. l Length is expressed in Angstrom (Å), nanometre (nm) and picometre (pm). These are related to the

SI unit as follows : 1 Å = 10–10 m, 1 nm = 10–9 m, 1 pm = 10–12 m. Ø Significant Figures—The digits in a properly recorded measurement or the total number of figures in a number

including the last digit whose value is uncertain are called significant figures, e.g., 180.00 has five significant figures. Ø Precision refers to the closeness of various measurements for the same quantity. Ø Accuracy is the agreement of a particular value to the true value of the results.Ø Laws of Chemical Combinations—Elements combined to form compounds in accordance to the following five

basic laws, called the laws of chemical combinations. (i) Law of Conservation of Mass—Mass can neither be created nor destroyed. (ii) Law of Definite Proportions—A given compound always contains exactly the same proportion of elements

by weight. (iii) Law of Multiple Proportions—When two elements combine with each other to form more than one

compound, the masses of one element that combine with a fixed mass of the other element, are in the ratio of small whole numbers.

(iv) Gay Lussac’s Law of Gaseous Volumes—When gases combine or are produced in a chemical reaction they do so in a simple ratio by volume provided all the gases are at same temperature and pressure.

(v) Avogadro’s Law—At the same temperature and pressure, equal volume of all gases contain the same number of molecules.

Ø Dalton’s Atomic Theory— l All matter consists of tiny indivisible particles. l Atoms are indestructible and unchangeable. l All atoms of a given element have identical properties, including identical mass. Atoms of different

elements differ in mass. l Elements are characterized by the mass of their atoms. l When elements react to form compounds, their atoms combine in simple whole number fixed ratio. l Chemical reactions involve reorganisation of atoms. These are neither created nor destroyed in a chemical

reaction.

Know the Terms Ø Matter : Anything that has mass and occupy space. Ø Accuracy : It refers to the agreements of a particular value to the true value of the results. Ø Precision : It refers to the closeness of various measurements for the same quantity. Ø Mass : Mass of a substance is the amount of matter present in the body. The mass of a substance is constant. Ø Weight : Weight is the force exerted by gravity on an object. The weight of a substance may vary from one place

to another due to change in gravity. Ø Scientific Notation : Expressing a number in the form N × 10n, and N can vary from 1.000... to 9.999... and n =

exponent having positive or negative values. Ø Significant Figures : These are meaningful digits which are known with certainty. Ø Mixture : A matter which is composed of two or more substances.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 3Know the Facts Ø The International System of Units (SI) : The International System of Units (in French Le System International

d’ Unites – abbreviated as SI) was established by the 11th General Conference on Weights and Measures (CGPM from Conference Generale des Poids et Measures). The SI system has seven base units.

Ø Volume : µ = 1 dm3 = 103 cm3 = 10–3 m3 .

Ø Temperature : K = °C + 273.15; °F - 32

9=

°C5

Ø Standard Temperature Pressure (STP) : 0 °C (273.15 K) temperature and 1 atmospheric pressure. Ø Normal Temperature Pressure (NTP) : 20 °C (293.15 K) temperature and 1 atmospheric pressure. ØStandard Ambient Temperature Pressure (SATP) : 25 °C (298.15 K) temperature and 1 atmospheric pressure.

TOPIC-2Different Types of Masses, Mole Concept and Formula of Compound

Quick Review Ø Atomic Mass : Relative mass of an atom is obtained by comparing its mass with 1/12th of the mass of one carbon

– 12 isotope atom.

Atomic Mass = 12

Mass of an atom(1/12) × Mass of a carbon atom (C )

Ø Average Atomic Mass : The average atomic mass of an element refers to the atomic masses of the isotopes of the element, taking into account the abundances of the element’s isotopes.

Average Atomic Mass =

S(Mass of isotopes × % natural abundance)

100

Ø Gram Atomic Mass : Gram atomic mass is the mass, in grams, of one mole of atoms in a monoatomic chemi-cal element.

Gram Atomic Mass = Atomic mass expressed in gram = Gram atom

Ø Molecular Mass : Molecular mass or molecular weight refers to the mass of a molecule. It is calculated as the sum of the mass of each constituent atom multiplied by the number of atoms of that element in the molecular formula.

Ø Gram Molecular Mass or Gram Molecule : Gram molecular mass is the mass in grams of one mole of a molecular substance.

Ø Formula Mass : The formula mass of a molecule is the sum of the atomic weights of the atoms in the em-pirical formula of the compound.

Ø Mole Concept and Molar Masses : Mole is a fundamental unit in the system. Molar mass is the mass of one mole of any substance i.e., element or compound.

Ø Percentage Composition : Percentage composition of a component in a compound is the percentage of the total mass of the compounds that is due to the component.

Ø Empirical Formula : A formula that gives the simplest whole number ratio of atoms in a compound.

Ø Molecular Formula : It is a formula in which the exact number of atoms of different elements are present.

Ø Some relations are follows :

Number of moles of atoms =

Mass of elementAtomic mass

Number of moles of molecules =

Mass of moleculesMolecular mass

Number of moles of gas =

Volume of the gas (STP)22·4 L

4 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Number of molecules = Number of moles × NA

Where NA = Avogadro’s number = 6·022 × 1023

Number of atoms = Number of molecules × Atomicity

Mass percentage =

×Total mass of element 100Molar mass

Know the Terms Ø Avogadro Number : It is the amount of atoms or molecules present in one mole of a substance.

Avogadro Number (NA) = 6.022 × 1023

Ø Mole (n) : It is the amount of a substance that contains as many particles or entites as the number of atoms is exactly 12 g of pure C-12.

Ø Molar Volume (Vm) : It is the volume occupied by one mole of any substance.

Ø Atomic Mass Unit (amu) : It is defined as a mass exactly equal to one-twelfth the mass of one C-12 atom.

Know the Facts Ø 1 amu = 1.66056 × 10–24 g

Ø Mass of hydrogen atom in terms of amu

=1 6736 10 24. ×

×

−

−g

1.66056 10 g/amu24

= 1.0078 amu Ø Now, ‘amu’ has been replaced by ‘u’ which is called unified mass. Ø Avogadro Number (NA) = 6.022 × 1023

Ø 1 mole of a substance = Molar mass of substance = Avogadro’s number of chemical units = 22.4 L volume at STP of gaseous substance

e.g. 1 mole of CH4 = 16 g of CH4 = 6.022 × 1023 molecules of CH4 = 22.4 L at STP

n = w

MV (at STP)

L particles

NMV1000

g

m

L

A= = =

22 4.x

Ø Molar volume of gas = 22.4 L at STP (273K, 1 atm) or 22.7 L at STP (273 K, 1 bar) Calculating Molar Volume : PV = nRT

∴ V = nRTP

=× ×− −1 mol 0.082 L atm K mol K

atm

1 1 2731

= 22.4 L

Or V = nRTP

=× ×− −1 mol 0.083 L bar K mol K

bar

1 1 2731

= 22.7 L

Ø Percentage Composition :

Mass % of the element = Mass of element in a molecule of the compound

Molecular masss of the compount×100

Ø Relationship between Empirical and Molecular Formulae: Molecular Formula = n × Empirical formula

where n = Molar mass

Empirical formula mass Ø Information Conveyed by a chemical equation:

N2(g) + 3H2 (g) → 2 NH3 (g)

(i) 1 molecule of N2 + 3 molecules of H2 → 2 molecules of NH3

(ii) 1 mole of N2 + 3 moles of H2 → 2 moles of NH3

(iii) 1 × 28 g of N2 + 3 × 2g of H2 → 2 × 17g of NH3

(iv) 1 × 22.4 L of N2 + 3 × 22.4 L of H2 → 2 × 22.4 L of NH3

at STP at STP at STP

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 5 Ø Relative abundance and Atomic mass of Isotopes of Carbon Isotope Relative Abundance Atomic Mass (%) (amu) 12C 98.892 12 13C 1.108 13.00335 14C 2 × 10–10 14.00317

TOPIC-3Stoichiometry and Reactions in Solutions

Quick Review Ø Chemical Equation—A chemical equation is the symbolic representation of a chemical reaction.

Ø Stoichiometry and Stoichiometric Calculations—

l Stoichiometry is the calculation of relative quantities of reactants and products in chemical reactions.

l Stoichiometric problems can be solved in just four simple steps : (i) Balance the equation. (ii) Convert units of a given substances to moles. (iii) Using the mole ratio, calculate the moles of substance yielded by the reaction. (iv) Convert moles of wanted substance to desired units.

Ø Mass-Mole-Number Relationship—

–1Mass of g

Number of Moles = Molar mass in g mol

Mass of

substance

Number of

ParticlesMultiplied by

molar mass

molar mass

Divided by Multiplied by

Avogadro's constant

Divided by

Avogadro'snumber

Mole

Ø Reactions in Solution—A solution consists of two or more substances dissolved in a liquid form.Ø Limiting Reactant—The reactant is entirely consumed when a reaction reaches completion. It is present in smaller

amount than calculated by balanced chemical equation.Ø Excess Reactant—The reactant which is not completely consumed in the reaction is called excess reactant.Ø Modes of expressing concentration of solution :

l Mass Percent—Mass percentage is one way of representing the concentration of an element in a compound or a component in a mixture.

×Mass of solute

Mass percent = 100Mass of solution

l Mole Fraction of component—The ratio of the number of moles of a given component of a mixture to the total number of moles of all the components.

Two components A and B Mole fraction of A

xA =

Number of moles of ANumber of moles of solution

A

A B= +

nn n

nA and nB are the mole of A and B respectively. Mole fraction of B,

xB =

Number of moles of B

Number of moles of solutionB

A B= +

nn n

Also remember that xA + xB = 1

6 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XIØ Molarity—A concentration unit, defined to be the number of moles of solute divided by the number of litres of

solution.

Ø Molality – The number of moles of solute per kilogram of solvent.

Molar mass of solute mass of solvent (ing)

Mass of solute 1000m =

Ø Strength of Solution—It is defined as the amount of solute dissolved per litre of solution. Strength = M × Molecular weight M = Molarity

Ø

% by strength =Volume of solute

Volume of solution× 100v

V

Ø

% by strength =

Weight of solute

Volume of solution× 100w

V

Ø

M =

% by mass × × 10Molecular weight of solute

d×100

Ø m = M

100 M Molecular weight× − ×d

d = density M = Molarity m = molalityØ Density =

MassVolume

Ø Specific gravity = Mass of liquidVol. of liquid

Ø Molecular weight = 2 × Vapour density

Ø For dilution, molarity equation is— M1V1 = M2V2

Where, M1 = initial molarity, V1 = initial volume, M2 = final molarity and V2 = final volume.

Know the Terms Ø Limiting Reagent : It is the reactant which gets consumed first or limits the amount of product formed. Ø Mass Percent : It is the mass of the solute in grams per 100 grams of the solution.

Mass % = Mass of solute in g ×100Mass of solution in g

Ø Parts per million (ppm). It is part of solute per million part of solution by mass.

ppm = Parts of solute (by mass) 10Parts of solution (by mass)

6×

Ø Mole Fraction (x) : It is the ratio of number of moles of one component to the total number of moles (solute and solvents) present in the solution.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 7Know the Facts Ø Molarity of a solution decreases on increasing temperature. Ø Molarity of pure water is 55.56 mol L–1. Ø Molarity is independent of temperature. Ø According to law of conservation of mass, a balanced chemical equation has the same number of atoms of each

element on both sides of the equation. Ø A balanced chemical equation provides a lot of information. Ø The quantitative study of the reactants required or the products formed is called stoichiometry.

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Chapter - 2 : Structure of Atom

TOPIC-1Dual Nature of Electromagnetic Radiation, Photoelectric Effect and Bohr Model

Quick Review Ø Developments Leading to the Bohr’s Model of Atom In order to improve Rutherford’s atomic model, two new concepts played a major role, these are : (i) Dual nature of electromagnetic radiation. (ii) Experimental results regarding atomic spectra. Ø Wave Nature of Electromagnetic Radiation—James Maxwell (1870) explained interaction between the charged

bodies and the behaviour of electrical and magnetic fields on microscope level. He suggested that when electrically charged particle moves under acceleration, alternating electrical and magnetic fields are produced and transmitted. These fields are transmitted in the form of waves called electromagnetic waves or electro magnetic radiation.

l These waves do not require any medium for their propagation and can move in vacuum. l The increasing order of wavelengths of these radiations is given as below : Cosmic rays < γ – rays < X – rays < Ultra–violet light < Visible light < Infra–red rays < Micro waves < Radio

waves. Ø Characteristics of Waves— l Wavelength (l)—The distance between two successive crests or troughs. l Frequency (u)—It is defined as the number of waves that pass a given point in one second.

l Wave number ( v )—The number of wavelengths per unit length. Ø Relationship between Wavelength, Wave number, Frequency and Velocity l These characteristics are related as :

c = l × u or u = cl and

= ul1 ∴ n = c v

Ø Particle Nature of Electromagnetic Radiation—The phenomenon like diffraction and interference can be explained by wave nature of electromagnetic radiation but it could not explain black body radiation and photoelectric effect.

Ø Black Body Radiation and its Explanation—The ideal body that emits and absorbs radiations of all frequencies, is called black body and the radiation emitted by such a body is called black body radiation.

l To explain the black body radiation, Planck gave his quantum theory. According to this theory, atoms and molecules emit (or absorb) energy in the form of small discrete packets of energy, known as quantum. The energy of each quantum is given by the expression :

E = hn where, h = Planck’s constant (6·626 × 10–34 J s) Ø Photoelectric Effect and its Explanation— l The photoelectric effect is the observation that certain metals (e.g., potassium, rubidium, caesium etc.) emit

electrons when light shines upon them. Electrons emitted in this manner can be called photoelectrons.

8 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Einstein explained the photoelectric effect using Planck’s quantum theory. l When a photon (packet of energy) with sufficient energy strikes an electron in the atom of the metal, it

transfers its energy instantaneously to the electron during the collision and that comes out without any time lag.

lGreater the energy possessed by the photon, greater will be transfer of energy to the electron and greater is the kinetic energy of the ejected electron. Kinectic energy of the ejected electron is proportional to the frequency of electromagnetic radiation.

l The minimum amount of energy required to eject the electron is known as wave-function and the extra amount of energy provided by the electron is used as kinetic energy of electron.

Total energy of photon = Work function + Kinetic energy.

hn = hn0 +

12

mv2

l For each metal, there is a characteristic minimum frequency known as threshold frequency (n0) below which photoelectric effect is not observed.

l Electromagnetic radiation shows wave like properties during propagation and shows particle like properties on interaction with matter.

Ø Dual Nature of Electromagnetic Radiation—Photoelectric effect shows particle nature of light. Diffraction and interference explain wave nature of light.

It shows that light possesses both particles and wave like properties, i.e., light has dual behaviour. Electromagnetic radiation shows wave like properties during propagation and shows particle like properties on

interaction with matter. Ø Evidence for the Quantized electronic energy levels : Atomic Spectra—The splitting of light into series of colour

bands is known as dispersion and the series of colour bands is called a spectrum. The branch of science that deals with the study of spectra is called spectroscopy.

Spectra are classified as: (i) Emission spectra (ii) Absorption spectra. Ø Emission Spectrum—The spectrum of radiation emitted by a substance that has absorbed energy is called an

emission spectrum. Ø The emission spectra of atoms in gas phase consist of bright lines against dark background as they emit only at

specific wavelengths. These spectra are called line emission spectra separated by bright bands. Ø The line spectrum of each element is unique and is therefore used for their identification in chemical analysis. Ø Absorption Spectra—When the white light is first passed through the substance and the transmitted light is

analysed we get absorption spectrum. It consists of dark lines separated by bright bands. Ø Line Spectrum of Hydrogen—All series of lines in the hydrogen spectrum can be described by Johannes

Rydberg expression :

u =

–12 21 2

1 1109, 677 – cmn n

Where n1 = 1, 2, ........ and n2 = n1 + 1, n1 + 2, ............. Ø The spectral Lines for Atomic Hydrogen are given as below : Series n1 n2 Spectral region Lyman 1 2, 3, ..... Ultraviolet Balmer 2 3, 4, ..... Visible Paschen 3 4, 5, ..... Infrared Brackett 4 5, 6, ..... Infrared Pfund 5 6, 7, ..... Infrared Ø From the line spectrum of different elements, it can be concluded that : (i) Line spectrum of an element is unique. (ii) There is a regularity in the line spectrum of each element. Ø Bohr’s Model for Hydrogen Atom In order to overcome the drawback of Rutherford’s model of atom and to explain the line spectrum of

hydrogen, Neils Bohr (1913) proposed a new model of atom based upon Planck’s quantum theory. Ø Postulates of Bohr’s Atomic Model— l In an atom, the electrons revolve around the nucleus in certain definite circular paths called orbits or shells. l Each shell or orbit corresponds to a definite energy. Therefore, these circular orbits are also known as energy

levels or energy shell. l The orbits or energy levels are characterized by an integer ‘n’, where, n can have values 1, 2, 3 or K, L, M ......... l The integral numbers are known as principal quantum number. l Electron can move from one orbital to another either by absorption of energy or by loss of energy.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 9 l The frequency of radiation absorbed or emitted is calculated by :

v

Eh

E Eh

= =−∆ 1 1

Where, E1 and E2 are the energies of the lower and higher allowed energy states respectively.

l Angular momentum of orbits in which electron revolves must be an integral multiple of p2

h , where h is Planck’s constant (h = 6.626 × 10–34 Js).

∴ mvr = n.p2

h

Where n = 1, 2, 3, .............. Ø According to Bohr’s Theory for Hydrogen Atom : (i) The radii of stationary states are given by : r n = n2a0 Where a0 = 52·9 pm and n = 1, 2, 3, .............. (ii) The energy of stationary state is given by :

En = – RH ( )21

n Where n = 1, 2, 3, ........ and RH (Rydberg constant) = 2·18 × 10–18 J. Ø For hydrogen like ions i.e., one electron system (e.g., He+, Li2+, Be3+), the radii is given by,

rn = 252·9( )

Zn pm

and energies of the stationary states are given by :

En = – 2·18 × 10–18 ( )2

2Zn

Ø Limitations of Bohr’s Model— l It fails to explain the doublet observed in the spectrum of hydrogen atom. l It fails to explain Zeeman effect (i.e., splitting of spectral lines in the presence of magnetic field) and stark

effect (i.e., splitting of spectral lines in the presence of an electric field). l It fails to explain the ability of atoms to form molecules by chemical bonds.

Know the Terms Ø Electromagnetic Radiation: Energy emitted from any source in the form of waves in which electric and magnetic

fields oscillated perpendicular to each other and travel with a velocity of light is called electromagnetic radiation. Ø Wavelength: The distance between one crest and one trough in a wave is called wavelength. It is denoted by the

symbol ‘l’. Ø Frequency: The number of waves passing through a given point in one second is called frequency. It is denoted

by the symbol ‘n’. Ø Wave Number: It is defined as the number of wavelength per unit length. It is denoted by the symbol n. Ø Amplitude: The height of crest or depth of a trough is called amplitude. It is denoted by the symbol ‘a’. Ø Velocity: The distance travelled by a wave in one second is called velocity. Ø Photon: A packet or particle of light energy is called photon. Ø Photoelectric Effect: The phenomenon of ejection of electrons from a metal surface when a light radiation of

suitable frequency falls on metal surface is called photoelectric effect. Ø Threshold Frequency: The minimum frequency (n0) below which photoelectric effect is not observed, is called

threshold frequency. Ø Line Spectrum: An emission spectrum is a photographic recording of the separated wavelengths is known as line

spectrum.

Know the Facts Ø The S.I. unit for frequency (n) is hertz (Hz, S–1), after Heinrich Hertz. Ø c = nl Where c = velocity of light (3 × 108 m/sec) n = frequency of light l = wavelength of light

10 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Planck gave the name quantum to the smallest quantity of energy that can be emitted or absorbed in the form of

electromagnetic radiation. Ø The energy (E) of a quantum of radiation is proportional to its frequency (n) and is expressed by equation: E = hn Where h = Planck’s constant (6.626 × 10–34 Joule-seconds) Ø Planck was awarded the Nobel Prize in Physics in 1918 for his quantum theory. Ø Albert Einstein, a German born American physicist, received the Nobel Prize in physics in 1921 for his explanation

of the photoelectric effect. Ø Energy wise order of electromagnetic radiation : Cosmic rays < γ-rays < X-rays < UV light < Visible light < IR rays < Micro waves < Radio waves

Wavelength (l) increases

Frequency (n) decreases

Wave number (v) decreases

Energy (e) decreases Ø In photoelectric effect, the difference in energy (hn – hn0) is transferred as the kinetic energy of the photoelectron.

hn – hn0 = 12

2mv

where hn = Light radiation falling upon metal surface hn0 = Energy used for work function or energy for removing electron from metal

12

2mv = Kinetic energy by which electron is emitted from metal surface.

Ø Large atoms have less work function. So, electron emits with more velocity Ø Small atoms have more work function. So, electron emits with less velocity. Ø Balmer showed in 1885 on the basis of experimental observations that if spectral lines are expressed in terms of

wave number (n), then the visible lines of the hydrogen spectrum obey the following formula :

vn

= −

109 6771

2

12 2, cm–1

Where n is an integer equal to or greater than 3 (i.e. n = 3, 4, 5 …) The series of lines described by this formula are called the Balmer series. Ø The value 109, 677 cm–1 is called the Rydberg constant for hydrogen.

TOPIC-2de-Broglie Equation, Heisenberg's Uncertainty Principle and Quantum Mechanical Model

Quick Review Ø Towards Quantum Mechanical Model of the Atoms—To overcome the shortcoming of the Bohr’s model, attempts

were made to develop a more suitable and general model for an atom. Two important developments which con-tributed significantly in the formulation of a new model were :

(i) Dual Nature of matter (ii) Heisenberg uncertainty principle. Ø Dual Behaviour of Matter—Einstein had suggested that light can behave as a wave as well as like a particle. i.e., it

has dual character. In 1924, de-Broglie suggested that matter and hence electron like radiations, has a dual character – wave and

particle. According to de-Broglie, ‘‘All material particles in motion possess wave characteristics’’. According to de-Broglie, the wavelength associated with a particle of mass m, moving with velocity n is given

by the relation,

λ = =hmv

hp

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 11 where h = Planck’s constant n = Velocity of particle (p = mv) = momentum of the particle. Ø Significance of de-Broglie equation Although the dual nature of matter is applicable to all materials but it is significant for microscopic bodies only. In case of ordinary objects, the wavelengths are so short that wave properties cannot be detected. Ø Heisenberg’s Uncertainty Principle Heisenberg’s uncertainty principle states that, “It is not possible to measure simultaneously both the position

and momentum (or velocity) of a microscopic particle, with absolute accuracy.” Mathematically, this law may be expressed as :

∆x × ∆p ≥

h4p

∆x.m∆v ≥

h4p

or ∆x × ∆v ≥

hm4p

where, ∆x = uncertainty in position, ∆p = uncertainty in momentum, ∆v = uncertainty in velocity. Ø Quantum Mechanical Model of Atom and Concept of Atomic Model— Quantum mechanics was developed independently by Heisenberg and Schrodinger.

Schrodinger equation is H^ ψ = E where H^ is a mathematical operator called Hamiltonian, E is the total energy

of the system and ψ is the wave function. Ø Important Features of the Quantum Mechanical Model of Atom 1. The energy of the electrons in atoms is quantized. 2. The existence of quantized electronic energy level is a direct result of the wave like properties of electrons. 3. Both the exact position and exact velocity of an electron in an atom cannot be determined simultaneously.

Therefore, only probability of finding an electron at different points is required. 4. An atomic orbital is the wave function ψ for an electron in an atom. 5. The probability of finding an electron at a point within an atom is proportional to the square of the orbital

wave function, i.e., |ψ|2 at that point. |ψ|2 is called probability density and is always positive. 6. From the value of |ψ|2 at different points within an atom, it is possible to predict the region around the

nucleus where the electron will most likely be found.

Know the Terms Ø Dual behaviour of matter : It means that matter behaves both as particle as well as wave nature. Ø Momentum : The quantity of motion of a moving body measured as a product of its mass and velocity. Ø Uncertainty in position of the particle : If the velocity of the electron is known precisely [∆(vx) is small], then the

position of the electron will be uncertain (∆x will be large), which is called uncertainty in position of the particle. Ø Uncertainty in velocity of the particle : If the position of the electron is known with high degree of accuracy (∆x

is small), then the velocity of the electron will be uncertain [∆(vx) is large], which is called uncertainty in velocity of the particle.

Ø Quantum mechanics : The branch of science that takes into account this dual behavior of matter is called quantum mechanics.

Ø Probability Density : The probability of finding an electron at a point with in an atom is proportional to the square of the orbital wave function i.e. |Y|2 at that point. |Y|2 is called probability density and is always positive.

Know the Facts Ø Electrons should have momentum as well as wave length. Ø Louis de Broglie was awarded the Nobel Prize in Physics in 1929. Ø The effect of Heisenberg uncertainty principle is significant only for the motion of microscopic objects and is

negligible for that of macroscopic objects. Ø Heisenberg was awarded the Nobel Prize in Physics in 1932.

12 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Quantum mechanics is a theoretical science that deals with the study of the motions of the microscopic objects

that have both observable wave like and particle like properties. It specifies the laws of motion that these objects obey.

Ø The fundamental equation of quantum mechanics was developed by Schrodinger and it won him the Nobel Prize in Physics in 1933.

Ø An atomic orbital is the wave function ψ for an electron in an atom.

TOPIC-3Atomic Orbital, Quantum Numbers and Electronic Configuration

Quick Review Ø Orbitals and Quantum Numbers

l Orbitals: By solving the Schrodinger wave equation ( ),H Eψ ψ= we obtain a set of mathematical equations, called wave functions (ψ), which describes the probability of finding electron in certain energy levels within an atom.

A wave function for an electron in an atom is called an atomic orbital, this atomic orbital describes a region of space in which there is a high probability of finding the electrons.

l Quantum Numbers: As discussed earlier, orbitals represent regions in space around the nucleus where the probability of finding the electrons is possible in an atom. These can be distinguished by their size, shape and orientation. Electron have only certain permissible values of energy and angular momentum. These permissible states of an electron in any atom are called orbitals which are identified by a set of four numbers. These numbers are called Quantum Numbers.

The various quantum numbers are : (a) Principal quantum number (n) (b) Azimuthal or angular momentum quantum number (l). (c) Magnetic quantum number (m l). (d) Spin quantum number (ms). (a) Principal quantum number (n) : This quantum number determines the energy associated with it. In addition,

it also determines the average distance of the electron from the nucleus in a particular shell. Starting from the nucleus, the energy shells are denoted as K, L, M, N, ..... etc. or as 1, 2, 3, 4, .... etc. The maximum number of electrons which a shell can accommodate is 2n2. Thus, K-shell (n = 1) can have a maximum of two electrons. L-shell (n = 2) can have eight electrons and similarly, eighteen electrons can be accommodated in M-shell (n = 3).

(b) Azimuthal or subsidiary or angular quantum number (l) : This quantum number determines the angular momentum of the electron. This is denoted by l. It identifies the subshell and determines the shape of the orbital. It can have positive integer values from zero to (n – 1) where n is the principal quantum number. i.e. l = 0, 1, 2, 3, ..... (n – 1).

Value of n

Value of l

Name of sub-shell

Number of the sub-shells

1 0 1s one2 0,1 2s, 2p two

3 0, 1, 2 3s, 3p, 3d three

4 0, 1, 2, 3 4s, 4p, 4d, 4f four

(c) Magnetic Quantum Number (ml): This describes the orientation of the orbitals along the three axis. Co-ordinate for a given value of l, the possible values of m vary from – l, 0 to + l.

(d) Spin Quantum Number (ms): This quantum number describes the spin orientation of the electron. It is designated by ‘(ms)’ Since the electron can spin in only two ways- clockwise or anticlockwise and, therefore, the spin quantum number can take only two values, + ½ or – ½.

Ø Shapes of Atomic Orbitals: It has been learnt that the probability of finding the electron does not become zero even at large distances from the nucleus. Therefore, it is not possible to draw any sort of geometrical figure that will enclose a region of 100% probability.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 13 Ø Radial Wave Function: An orbital is a mathematical function called a wave function that describes an electron in

an atom.

The wave function (ψ), of the atomic orbitals can be expressed as the product of a radial wave function.

where ψ = Wave function, r = Distance

Ø Probability Density (ψ2) Graphs: German physicist, Max Born suggested that the probability density of the electron at any point is obtained by the square of the wave function ψ2 at that point.

These graphs are known as probability density graphs.

Ø Boundary Surface Diagrams: The shapes of atomic orbitals may be represented in terms of boundary surface diagram.

These boundary surface diagrams of some atomic orbitals are discussed below :

14 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI l Shapes of s-orbitals: s-orbitals are non-directional and spherically symmetrical.

l Shapes of p-orbitals: For p-orbitals (l = 1), there are three possible orientations corresponding to ml = – 1, 0, + 1 values. This means that there are three p-orbitals in each p-sub-shell.

These are designated as px, py, and pz.Z

Y 2 -orbitalpx

(a)

Z

Y 2 -orbitalpy

(b)

X X

Z

Y2 -orbitalp

z

(c)

X

l Shapes of d-orbitals: The five-d-orbitals, designated as dxy, dyz, dzx, dx2-y2, dz2. The shapes of first four d-orbitals are similar to each other, whereas fifth one dz2 is different from others but all the five d-orbitals are equivalent in energy i.e., degenerate.

Z

Y

dxy

X

Z

Y

X

dxz

Z

YX

dx y2 – 2

Z

YX

dyz

Z

YX

dz2

Fig. Shapes of five d-orbitals

Ø Nodes and Nodal Planes: The place or point where radial wave function passed through zero are called nodes or radial nodes.

This point at which radial wave function becomes zero is called radial nodal surface. Ø Energies of Orbitals: Energies of different orbitals of hydrogen and hydrogen like particles depend only upon

the value of principal quantum number (n). But in case of multielectrons atom, it depends on the value of principal quantum number as well as azimuthal

quantum number. Ø Degenerate Orbitals: The orbitals with same energy are called degenerate orbitals e.g., p sub-shell has three

degenerate orbitals namely px, py and pz. Ø Shielding Effect: The term shielding effect refers to a decrease in attraction between electrons and the nucleus

in an atom. Electrons are highly attracted to the nucleus, because they have a negative charge and the nucleus contains protons, which have a positive charge.

Ø n + l Rule: Orbitals are filled in the order of increasing n + l. For equal n + l values, the orbital with the lower n is most often filled first. Here n is the principal quantum number and l is the angular momentum quantum number.

Ø Filling of Electrons in Atomic Orbitals: The filling of electrons in different orbitals is governed by the following rules : (i) Aufbau Principle: Electrons are filled in the various orbitals in an increasing order of their energies, i.e.,

orbital having lowest energy will be filled first and the orbital having highest energy will be filled last.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 15

Fig. Order of filling of electrons in orbitals

Ø Pauli’s Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers. It can also be stated as ‘an orbital can have maximum two electrons and they must be of opposite spin quantum numbers’.

e.g., two electrons in an orbital can be indicated either by (i) or (ii) but, not by (iii) or (iv)

Correct (i) or (ii)

Incorrect (iii) or (iv)

Ø Hund’s Rule of Maximum Multiplicity: No electron pairing takes place in p, d and f orbitals until each orbital in the given sub-shell contains one electron having parallel spin, e.g., N (7) has electronic configuration 1s2, 2s2, 2px1, 2py1, 2pz1 according to Hund’s rule and not 1s2, 2s2 2px2 2py1.

1s 2s 2px 2py 2pz

Ø Electronic Configuration of Atoms: The filling of orbitals in an atom is a hypothetical process in which the atom is built up by feeding electrons in orbitals, one at a time and by placing each new electron in the lowest available energy orbital.

The distribution of electrons in different orbitals is known as electronic configuration of the atom. Ø Significance of Electronic Configuration: The number of valence electrons in an atom has determined how

readily an atom will react with other atoms. If it has 7 or 1 valence electron, then it only needs to gain/loss 1 electron which is relatively easier than gaining/losing 3 or 4 electrons.

Ø Electronic Configuration of Ions: The electronic configuration of the ions may be written almost similar to those of the atoms. In case of negatively charged ions, the extra electrons equal to the charge are added to the appropriate orbitals.

Examples :

Cl (No. of electrons = 17) : 1s2, 2s2, 2p6, 3s2, 3px2 , 3py

2 , 3pz

1

Cl– (No. of electrons = 17) : 1s2, 2s2, 2p6, 3s2, 3px2 , 3py

2 , 3pz

2

Mg (No. of electrons = 12) : 1s2, 2s2, 2p6, 3s2

Mg2+ (No. of electrons = 12) : 1s2, 2s2, 2p6

16 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI

Know the Terms Ø Nodal surface: The region where the probability density function reduces to zero is called nodal surfaces or

simply nodes. Ø Degenerate orbitals : The orbitals having the same energy are called degenerate orbitals. Ø Effective nuclear charge: It is the net charge an electron experiences in an atom with multiple electrons. It may

be approximated by the equation Zeff = Z – S where Z = atomic number S = number of shielding electrons Ø Electronic configuration : The distribution of electrons into orbitals of an atom is known as its electronic

configuration.

Know the Facts Ø An orbital of small size means there is more chance of finding the electron near the nucleus. Ø An orbit is a circular path around the nucleus in which an electron moves. However, an atomic orbital is a quantum

mechanical concept and refers to the one electron wave function ψ in an atom. Ø The total probability of finding the electron in a given volume can be calculated by the sum of all the products of

|ψ|2 and the corresponding volume elements. Ø The orbital wave function or ψ for an electron in an atom has no physical meaning. It is simply a mathematical

function of the coordinates of the electron. Ø According to the German physicist, Max Born, the square of the wave function (i.e. ψ2) at a point gives the

probability density of the electron at that point. Ø If probability density |ψ|2 is constant on a given surface. |ψ| is also constant over the surface. The boundary

surface for |ψ|2 and |ψ| are identical.

qq

Chapter - 3 : Classification of Elements and Periodicity in Properties

TOPIC-1Classification of Elements and Periodic Table

Quick Review Ø Need of Classification—Upto the end of seventeenth century, only 31 elements were known. However, at present

114 elements are known. Therefore, it is not easy to study and remember the properties of these elements. By arranging the elements in such a way that similar elements are placed together while dissimilar elements are

separated from one another. This is known as classification of elements. Such a classification of the elements has resulted in the formulation of the periodic table. Ø Periodic table may be defined as : The arrangement of the known elements according to their properties in a tabular form. Ø Historical Development of Periodic Table— Since the beginning of the ninteenth century, scientists have been trying to find a basis of grouping elements

having similar properties. Some of the earlier important attempts to classify the elements are briefly summed up below :— (i) Dobereiner’s triads—In 1817, a German scientist, Johann Dobereiner classified the elements in groups of

three elements called triads. The elements in a triad had similar properties and the atomic weight of the middle member of each triad is

very close to the arithmetic mean of the other two elements.DOBEREINER’S TRIADS

Li Lithium Ca Calcium Cl Chlorine

Na Sodium Sr Strontium Br Bromine

K Potassium Ba Barium I Iodine

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 17 Elements Atomic weight

Lithium 7 Sodium 23 Potassium 39 Atomic weight of Na = 23

Mean of other two elements = 7 39

2+

= 23 (ii) A.E.B de Chancourtois Classification – He arranged the known elements in order of increasing atomic

weights and proposed a cylindrical table of elements to display the periodic recurrence of properties. (iii) Newlands’ Law of Octaves–When the elements are arranged in the increasing order of atomic weights, the

properties of every eighth element are similar to the first one. Newlands called this relation as the law of octaves due to similarity with the musical scale. (iv) Lothar Meyer’s Arrangements – When the properties of the elements such as atomic volume, melting point,

boiling point etc., are plotted as a function of their atomic weights, the elements with similar properties occupied almost similar positions.

(v) Mendeleev’s Periodic Law – The physical and chemical properties of elements are periodic function of their atomic weights.

Ø Characteristics of Mendeleev’s Periodic Table— It consists of (i) Eight vertical columns, called groups—Except for VIII group, each group is further sub-divided in A and B.

This sub division is made on the basis of difference in their properties. (ii) Six horizontal rows, called periods. Significance of Mendeleev’s Periodic Table— (i) Instead of studying properties of elements separately, they can be studied in groups containing elements

with same properties. It led to the systematic study of the elements. (ii) Prediction of new elements : At his time only 56 elements were known. He left blank spaces for unknown

elements. (iii) Mendeleev’s periodic table corrected the doubtful atomic weights. Ø Defects in the Mendeleev’s Periodic Table : (i) Hydrogen was placed in group IA—However, it resembles both group IA elements (alkali metals) and VII A

(halogens). Therefore the position of hydrogen in the periodic table is anomalous or controversial. (ii) Anomalous pairs of elements—Some elements with higher atomic weight like Argon (39.9) precede

potassium (39.1) with lower atomic weight. (iii) Based upon atomic weights, isotopes of an element could be assigned different groups but as they are the

atoms of same element (with only difference in their molecular mass) they must be in same group. (iv) Some dissimilar elements are grouped together while some similar elements are placed in different

groups. For example, alkali metals in group IA which are highly reactive are in the same group as coinage metals like Cu, Ag, Au of group IB. At the same time, certain chemically similar elements like Cu (group IB) and Hg (group IIB) have been placed in different groups.

(v) Position of elements of group VIII—No proper place has been allotted to nine elements of group VIII which have been arranged in three triads without any justification.

Ø Modern Periodic Law : The physical and chemical properties of the elements are periodic functions of their atomic number.

According to the recommendations of IUPAC, the groups are numbered from 1 to 18 replacing the older notation of groups 0, IA, II A ......

There are 7 periods. The first period contains 2 elements. The subsequent periods contain 8, 8, 18, 18, 32 elements respectively. The 7th period is incomplete and the 6th period would have a theoretical maximum of 32 elements. In this form of periodic table, the elements of both sixth and seventh periods (lanthanoids and actinoids respectively) are placed in separate panels at the bottom.

Ø Long Form of Periodic Table : General characteristics of the long form of the periodic table— (i) There are in all, 18 vertical columns or 18 groups in the long form periodic table. (ii) These groups are numbered from 1 to 18 starting from the left. (iii) There are seven horizontal rows called periods in the long form of periodic table. Thus, there are seven

periods in the long form of periodic table. First period 2 elements Shortest period Second and Third period 8 elements each Short period Fourth and Fifth period 18 elements each Long period Sixth period 32 elements Longest period Seventh period It is incomplete Incomplete period


(iv) The elements of group 1, 2 and 13 to 17 are called the main group elements. These are also called typical or representative or normal elements.

(v) The elements of group 3 to 12 are called transition elements. (vi) Elements with atomic number 58 to 71 (Ce to Lu) occurring after lanthanum (La) are called Lanthanoids.

Elements with atomic number 90 to 103 (Th to Lr) are called Actinoids. These elements are called f-block elements and also as inner-transition elements.

Ø Defects of long form of the periodic table : l In periodic table, the position of hydrogen is also not clear. l It is placed in the separate box just above the group I. l Lanthanoids and actinoids have not been merged within the main body of the periodic table.

Know the Terms Ø Periods : The horizontal rows in the periodic table are called periods. Ø Groups : The vertical columns in the periodic table are called groups.

Know the Facts Ø In 1800, only 31 elements were known. Ø By 1865, the number of identified elements had more than doubled to 63. Ø At present, 114 elements are known. Ø Johann Dobereiner (1800) was the first to consider the idea of trends among properties of elements. Ø In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. Element 106 has been named Seaborgium

(Sg) in his honour. 1st period — 2 elements. 2nd and 3rd period — 8 elements. 4th and 5th period — 18 elements. 6th period — 32 elements. 7th period — incomplete. Ø In modern periodic table, there are 18 groups. Group 1 and 2 : s-block elements in which last electron entered in s-sub shell [s1, s2] Group 3 to 12 : d-block elements in which last electron entered in d-sub shell [d1 to d10] Group 13 to 18 : p-block elements in which last electron entered in p-sub shell [p1 to p6] Group 18 : Noble gases. Ø In ‘s’ and ‘p’ – block elements, the last electron enters in the outermost shell. Ø In d-block elements, the last electron enters in the penultimate shell (n – 1). Ø In f-block elements, the last electron enters in the sub-penultimate shell (n – 2).

TOPIC-2Nomenclature and Relation between Electronic Configuration and Periodic Table

Quick Review Ø Nomenclature of Elements with Atomic Number Greater than 100— It is a generally accepted convention that the discovery of an element has the honour of naming it. (i) The name is derived directly from the atomic number of the element using the latin numeral roots. (ii) The symbol of the element is composed of the initial letters of the numeral roots which make up the name. (iii) The root ‘un’ is pronounced with a long ‘u’ to rhyme with ‘moon’. In the element names, each root to be

pronounced separately. Ø Electronic Configuration of Elements and the Periodic Table— For the sake of simplicity, the electronic configuration of the noble gas elements are written as [He]2, [Ne]10, [Ar]18,

[Kr]36, [Xe]54 and [Rn]86. The Aufbau (build up) principle and the electronic configuration of atoms provide a theoretical foundation for

the periodic classification. The elements in a vertical column of the Peridic Table constitue a group or family and

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 19exhibit similar chemical behaviour. This similarity arises because these elements have the same number and same distribution of electrons in their outermost orbitals. We can classify the elements into four blocks viz, s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons.

Ø Electronic Configuration in Periods—The period indicates the value of n for the outermost or valence shell. In other words, successive period in the Periodic Table is associated with the filling of the next higher principal energy level (n = 1, n = 2 etc.). It can be readily seen that the number of elements in each period is twice the number of atomic orbitals available in the energy level that is being filled.

(i) Elements in first period—The first period (n = 1) starts with the filling of the lowest level (1s) and therefore has two elements – hydrogen (1s1) and helium (1s2) when the first shell (K) is completed.

(ii) Elements in second period—The second period (n = 2) starts with lithium and the third electron and has the electronic configuration 1s22s1. Starting from the next element boron, the 2p orbitals are filled with electrons when the L shell is complete at neon (2s22p6). Thus, there are 8 elements in the second period.

(iii) Elements in third period—The third period (n = 3) begins at sodium, and the added electrons enter a 3s orbital. Successive filling of 3s and 3p orbitals give rise to the third period having 8 elements from sodium to argon.

(iv) Elements in fourth period—The fourth period (n = 4) starts at potassium, and the added electrons fill up the 4s orbital. Also, 3d orbitals are filled before 4p orbitals because 3d has lower energy. Fourth period ending at krypton has 18 elements.

(v) Elements in fifth period—It starts with Rb (37) with 5s orbital followed by filling of 4d orbitals, then 5p orbitals and ends at Xe (54), having total 18 elements.

(vi) Elements in sixth period—It has 32 elements and successive electrons are filled in 6s, 4f, 5d and 6p orbitals, in that order.

(vii) Elements in seventh period—This period corresponds to n = 7 just like sixth period it involves the filling of 7s, 5f, 6d and 7p. This period contains synthetic or man-made radioactive elements.

Ø Groupwise Electronic Configuration–Elements in a group have similar electronic configuration, same number of electrons in valence orbitals and similar properties e.g., Group 1 (alkali) metals have general electronic configuration ns1.

Ø Position of Helium—The electronic configuration of helium is 1s2. So it should be in s-block, but it is placed in p-block due to the presence of completely filled valence shell (1s2) and similarity in properties with other noble gases.

Ø Position of Hydrogen—Electronic configuration of hydrogen is 1s1. Because of the presence of only one electron in s-orbital, it can be placed in group 1 with alkali metals and it also achieves the electronic configuration of inert gas so, it can also be placed in 17 group.

Ø s-Block Elements—The elements of group 1 (alkali metals) and group 2 (alkaline earth metals) which have ns1 and ns2 outermost electronic configuration belong to the s-block elements. They are all reactive metals with low ionization enthalpies. They lose the outermost electron(s) readily to form +1 ion (in the case of alkali metals) or +2 ion (in the case of alkaline earth metals). The metallic character and the reactivity increase as we go down the group. Because of high reactivity they are never found pure in nature. The compounds of the s-block elements, with the exception of those of lithium and beryllium are predominantly ionic.

Ø p-Block Elements—The p-block elements comprise those belonging to groups 13 to 18 and these together with the s-block elements are called the Representative Elements or Main Group Elements. The outermost electronic configuration varies from ns2np1 to ns2np6 in each period. At the end of each period is a noble gas element with a closed valence shell ns2np6 configuration. All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity preceding the noble gas family are chemically important groups of non-metals. They are the halogens. The non-metallic character increases as we move from left to right across a period and metallic character increases as we go down the group.

Ø The d-Block Elements (Transition Elements)—These are the elements of Group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner (n – 1) d orbitals by electrons and are therefore referred to as d-block Elements. These elements have the outer electronic configuration (n – 1)d1-10ns1-2. They are all metals.

However, Zn, Cd and Hg which have the electronic configuration, (n – 1)d10 ns2 do not show most of the properties of transition elements. In a way, transition metals form a bridge between the chemically active metals of s-block elements and the less active metals of Groups 13 and 14 and thus take their familiar name “transition elements”.

Ø The f-Block Elements (Inner-Transition Elements)—The two rows of elements at the bottom of the Periodic Table, called Lanthanoids, Ce (Z = 58) – Lu (Z = 71) and Actinoids, Th (Z = 90) – Lr (Z = 103) are characterised by the outer electronic configuration (n – 2)f1-14 (n – 1 )d0-1 ns2. The last electron added to each element is f-electron. These two series of elements are hence called the inner transition elements (f-Block Elements).

They all are metals. The chemistry of the early actinoids is more complicated than the corresponding lanthanoids due to large number of oxidation states.


Division of Periodic Table into s, p, d, f-Blocks

Ø Metals, Non-metals and Metalloids—The elements can be divided into Metals and Non-metals. Ø Metals— l Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table. l Metals are usually solid at room temperature (Mercury is an exception). l Metals usually have high melting and boiling point except gallium and caesium. Gallium and caesium have

very low melting point 303 K and 302 K). Ø Non-Metals—These are present only in p-block and appear on the right hand side of the Periodic Table. l These can be either solid or gases at room temperature. (exceptional case is bromine which is liquid at room

temperature). l Melting and boiling points are very low except boron and carbon. l They are brittle and neither malleable nor ductile. Ø Metalloids—The elements like silicon, germanium, arsenic, antimony and tellurium (Si, Ge, As, Sb, Te) which

show properties of both metals and non–metals are called Metalloids.

Know the Terms Ø s-Block Elements : The elements in which the last electron enters in the s-orbital of their respective outermost

shell are known as s-block elements. Ø p-Block Elements : The elements in which the last electron enters in the p-orbital of their respective outermost

shell are known as p-block elements. Ø d-Block Elements : Elements in which the last electron enters in one of the inner d-orbital are known as d-block

elements. There are also called transition elements. Ø f-Block Elements : Elements in which the last electron enters in f-orbital of anti-penultimate shells are known as

f-block elements. These are also called inner-transition elements. Ø Lanthanoids : The elements of 4f series [from atomic number 58(Ce) to 71(Lu)] are called lanthanoids. Ø Actinoids : The elements of 5f series [from atomic number 90 (Th) to 103 (Lr) are called actinoids. Ø Transuranic Elements : The elements which follow uranium [i.e. neptunium to Uub (z = 112) of actinoids] in the

periodic table are called transuranic elements. Ø Metalloids : The elements which exhibit the properties of both metals and non-metals are called metalloids.

Know the Facts Ø The period indicates the value of n for the outermost or valence shell. Ø Succcessive period in the periodic table is associated with the filling of the next higher principal energy level (n =

1, n = 2 etc) Ø The distribution of electrons into the orbitals of an atom is called its electronic configuration. Ø All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to

alter this stable arrangement by the addition or removal of electrons. Ø Actinoid elements are radioactive.

TOPIC-3Periodic Properties and their Periodic Trends

Quick ReviewØ Atomic Radii : The term atomic radius means the distance from the centre of the nucleus to the outermost shell

of electrons.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 21Ø On the basis of the nature of combining atoms, atomic radius can be of following types : l Covalent Radius : Covalent radius may be defined as “one half of the distance between the centre of nuclei

of two similar atoms bonded by a single covalent bond. For homonuclear molecules,

=covalentInternuclear distance between two bonded atoms

2r

l Metallic Radius : Metallic radius is taken as-one half of the internuclear distance between the two neighbouring atoms of a metal in a metallic lattice.

l van der Waals’ Radius : Noble gases do not have molecules, so it is not possible to measure covalent or ionic radius. Their atoms are held together by weak van der Waals forces of attraction. The van der Waals radius is half of the distance between the centre of nuclei of atoms. Noble gases held together by weak van der Waals forces of attraction.

Ø Variation of Atomic Radii in the Periodic Table : The following periodic trends in atomic radii have been observed: l Variation in a period : In general, the atomic radii decrease with increase in atomic number on going from

left to right due to the increase of effective nuclear charge. For example, in the second period, the atomic radii decrease from Li to Ne through Be, B, C, N, O and F. l Variaton in a group : The atomic radii of elements increase from top to bottom in a group. On moving down

a group, the effective nuclear charge decreases with increase in atomic number and we expect that the size of atom should increase.

Ø Ionic Radii : The effective distance from the centre of the nucleus of the ion upto which it has an influence in the ionic bond.

The ionic radii show the same trends as atomic radius.

The study of ionic radii leads to two important generalizations :

(i) The radius of positive ion (cation) is always smaller than that of the parent atom because it has lesser electrons but same nuclear charge.

(ii) The radius of negative ion (anion) is larger than that of the parent atom because extra electrons are added due to which repulsion among electrons increases and effective nuclear charge decrease.

Ø Isoelectronic Species : These are the atoms or ions having the same number of electrons. e.g., O2–, F–, Na+ and Mg2+ have 10 electrons. Among these species, the cation with greater positive charge will have a smaller radius because of the greater attraction of the electrons to the nucleus. However, the anion with the greater negative charge will have the larger radius as the net repulsion of the electrons outweigh the nuclear charge.

Ø Ionization Enthalpy : It is defined as the energy required to remove an electron from isolated gaseous atom in its ground state.

The first ionistion enthalpy (IE1) is the energy required to remove the most loosely bound electron of the neutral atom and the second ionisation enthalpy, IE2 is the energy required to remove the second electron from the resulting cation and so on. Thus first ionisation enthalpy of an element (X) may be defined as reaction represented as :

X(g) →1IE

X+(g) + e–(g)

Ø Factors Affecting Ionization Enthalpy : The magnitude of ionisation enthalpy for an atom depends upon the following factors :

l Size of the atom : The ionisation enthalpy depends upon the distance between the electron and the nucleus, i.e., size of the atom.

l Charge on the nucleus : The attractive force between the nucleus and the electron increases with the increase in nuclear charge. This is because, the force of attraction is directly proportional to the product of charges on the nucleus and the electron.

l Screening effect : The reduction in force of attraction by the electrons of shells present in between the nucleus and valence electrons is called screening effect or shielding effect. The greater the number of intervening electrons between valence electron and nucleus, the greater will be shielding or screening effect.

l Penetration effect of electrons : s-electrons are more penetrating towards the nucleus than p-electrons and the penetration power decreases in a given shell in the order.

s > p > d > f

For the same shell, the ionization enthalpy would be more to remove an s-electron than the energy required to remove a p-electron, which in turn will be more than that for the removal of a d-electron and so-on.

22 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Electronic Arrangement : The ionization enthalpy also depends upon the electronic configuration of the atom. For example, half filled and completely filled shells have extra stability associated with them. This may be illustrated by the following examples : l The noble gases have the most stable electronic configurations (ns2np6) in each period and therefore, have

high ionization enthalpies. l The elements like Be (1s2 2s2) and Mg (1s2 2s2 2p6 3s2) have completely filled orbitals and their ionisation

enthalpies are large. Ø Variation Down a Group : Within a group, there is a gradual decrease in ionization enthalpy in moving from top

to bottom. The decrease in ionization enthalpy down a group can be explained in terms of net effect of the following factors: l While going from top to bottom in a group, the nuclear charge increases. l There is a gradual increase in atomic size due to an additional main energy shell (n). l There is increase in shielding effect on the outermost electron due to increase in the number of inner electrons

that outweighs the increasing nuclear charge and removal of the outermost electron requires less energy down a group.

Ø Variation along a period : In general the ionisation enthalpy increases with increasing atomic number in a period.

Variation of ionization enthalpy with atomic number in second period

The general increase of ionization enthalpy along a period can be explained on the basis of atomic size and effective nuclear charge.

We know that : (i) On moving across a period from left to right, the effective nuclear charge increases. (ii) The atomic size decreases along a period though the main energy level remains the same. Ø Successive Ionization Enthalpies : The energies required to remove subsequent electrons from the atom in the

gaseous state are known as successive ionization enthalpies. It has been found that

IE1 < IE2 < IE3 Ø Electron Gain Enthalpy (∆egH) : It is defined as the enthalpy change when a neutral gaseous atom takes up extra

electrons to form an anion.X(g) + e– —→ X–(g)

The value of electron gain enthalpy depends upon atomic size, nuclear charge etc. With increase in size, the electron gain enthalpy decreases as the nuclear attraction decreases. Thus electron gain enthalpy generally decreases in moving from top to bottom in a group (F < Cl > Br > I). In a period from left to right, the electron gain enthalpy generally increases due to increase in effective nuclear charge.

Ø Factors Affecting Electron Gain Enthalpy : There are many factors which govern the electron gain enthalpy but the following are some important factors on which it mostly depends :

l Nuclear charge : The electron gain enthalpy becomes more negative as the nuclear charge increases. This is due to greater attraction for the incoming electron if nuclear charge is high.

l Size of the atom : With the increase in size of the atom, the distance between the nucleus and the incoming electron increases and this results in lesser attraction.

Consequently, the electron gain enthalpy becomes less negative with increase in size of the atom of the element.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 23 l Electronic configuration : The elements having stable electronic configuration of half filled and completely

filled valence sub-shells show very small tendency to accept additional electron and thus electron gain enthalpies are less negative.

Ø Important Trends in Electron Gain Enthalpy : There are some important features of electron gain enthalpies of elements. These are :

l Halogens have the highest negative electron gain enthalpies : The electron gain enthalpies of the halogens are highly negative. This is due to the fact that halogens have the general electronic configuration of ns2np5 and have only one electron less than the stable noble gas (ns2np6) configuration.

l Electron gain enthalpy values of noble gases are positive while those of Be, Mg, N and P are almost zero : The electron gain enthalpy values of noble gases are positive. This is because they have stable electronic configuration of ns2np6 and thus they have absolutely no tendency to take an additional electron. This means that the incoming electron enters the next higher principal quantum level and does not feel any attraction for the nucleus. Thus, energy is required to force the electron in their atoms and therefore, their electron gain enthalpies are positive.

l Electron gain enthalpy of fluorine is unexpectedly less negative than of chlorine : As already explained, the electron gain enthalpy, in general, becomes less negative from top to bottom in a group. However, it is observed that F-atom has unexpectedly less negative electron gain enthalpy than Cl-atom. This is due to the very small size of F-atom. As a consequence of small size, there are strong inter electronic repulsions in the relatively compact 2p-subshell of fluorine and thus the incoming electron does not feel much attraction.

Ø Electronegativity : l It is a qualitative measure of the ability of an atom in a chemical compound to attract shared pair of electrons

to itself. l It is not a measurable quantity, however the electronegativity of different elements on different scales was

developed. Pauling scale is generally used on which fluorine (most electronegative element) was assigned a value of 4.0 arbitrarily.

l Electronegativity generally increases on moving across a period because of increase in effective nuclear charge.

l It decreases on moving down a group because of decrease in effective nuclear charge. Ø Summary of the trends in the periodic properties of elements in the periodic table :

Ø Valency : The electrons present in the outermost shell of an atom are called valence electrons and the number of these electrons determine the valence or the valency of the atom. It is because of this reason that the outermost shell is also called the valence shell of the atom and the orbitals present in the valence shell are called valence orbitals.

Ø Diagonal Relationship—

The resemblance of properties of Li with Mg (which is diagonally situated) : Be with Al and B with Si is called diagonal relationship.

It is due to similar polarising power.

24 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Periodic Trends and Chemical Reactivity— l As we have seen, the ionization enthalpy of the extreme left element in a period is the least and the electron

gain enthalpy of the element on the extreme right is the highest negative (excluding noble gases). This results in high chemical reactivity at the two extremes and the lowest in the centre.

l Thus, alkali metals (at extreme left) has maximum tendency to form cation (by loss of an electron) and halogens (at extreme right) has maximum tendency to form anion (by gain of an electron).

l Thus, metallic character is maximum at the left and non-metallic character is maximum at the right within a period.

l The oxides formed by the elements at the left of the periodic table are basic (e.g., Na2O) while oxides formed by the elements at the right of the periodic table are acidic (e.g., Cl2O7). However, oxides of elements in the centre are either amphoteric (e.g., Al2O3, Sb2O3) or neutral (e.g., CO, NO, N2O).

l In transition metals, change in atomic radii is much smaller than representative elements. Their ionization enthalpies are between s-and p-block elements. Therefore, they are less electropositive than group 1 and 2 metals.

Know the Terms Ø Atomic Radius : It is defined as the distance between the centre of the nucleus and outermost shell containing

electrons.

Ø Covalent Radius : It is defined as the one-half of the distance between the centres of nucleus of two covalently bonded atoms of the same element in a molecule.

Ø Metallic Radius : It is defined as the one-half of the inter-nuclear distance separating the metal cores in the metallic crystal.

Ø Van der Waals’ Radius : It is defined as the one-half of the distance between the nuclei of two identical non-bonded isolated atoms.

Ø Shielding Effect : The electrons of inner shells (core electrons) shield or screen the outer most electrons from the attractive pull of the nucleus. Such type of effect is called shielding effect. It is also called screening effect.

Ø Effective Nuclear Charge-(Zeff) : It is defined as the net nuclear attraction towards the valence shell electrons.

Ø Ionic Radius : It is defined as the effective distance from the centre of the nucleus of an ion upto which it has an influence in the ionic bond.

Ø Ionization Enthalpy : The energy required to remove an electron from an isolated gaseous atom in its ground state is known as ionization enthalpy.

Ø Electron gain Enthalpy : When an electron is added to a neutal gaseous atom (X) to convert it into a negative ion, the enthalpy change accompanying the process is known as electron gain enthalpy.

Ø Electronegativity : A quantitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is known as electronegativity.

Know the Facts Ø General Trends of Periodic Properties in Periodic Table (a) Trends of Atomic Radius (i) Decreases on moving from left to right in a period due to the effect of successive increasing nuclear charge

without addition of a new shell. (ii) Increases from top to bottom in a group due to successive addition of shell. (iii) Noble gases have large atomic radius than group 17 elements due to complete filling of electron in outer shell

electron-electron repulsion mildly increases. (b) Trends of Ionization Enthalpy (i) Increases on moving from left to right in a period due to decrease in atomic radius and for 2nd period due to more

force of attraction by protons First Ionization Enthalpy : Li < B < Be < C < O < N < F <Ar Second Ionization Enthalpy : Be <C , B < N < F < O < Ne (ii) Decreases from top to bottom in a group due to increase in atomic size or due to addition of shells.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 25 Ø Metallic Behaviour : Decreases from left to right in a period due to increase in ionisaion enthalpy.

Ø Non-metallic Behaviour : Increases from left to right due to more number of electron in outermost shell and added electron goes towards nucleus.

Ø Radius of Cation< Atomic radius It is due to more number of protons than the number of electrons in a cation. Coulombic force increases, size

decreases. Mg2+< Mg+< Mg

Ø Radius of Anion > Atomic Radius It is due to more number of electrons than the number of protons in an anion. N3– > O2– > F–

Electron – electron repulsion increases, coulombic force of attraction decreases and size increases.

Ø For Isoelectronic Species : l Greater the charge of cation, lesser will be the size. l Greater the charge of anion, greater will be the size. Eg. O2– > F– > Na > Na+ > Mg2+

Ø Trends of Electron Gain Enthalpy : (a) Increases from left to right in a period due to decrease in size and more attraction of added electron by nucleus. (b) Decreases from top to bottom in a group as the added electron is away from nucleus due to increase in size. Halogen > Oxygen > Nitrogen > Metal of group 1 and 13 and non-metal of group 14 > Metal of group 2. (c) Chlorine has maximum electron gain enthalpy. Cl has more electron gain enthalpy than F due to small size of

fluorine, extra added electron has more inter electronic repulsion than chlorine. (d) Phosphorus and sulphur have negative electron gain enthalpy then nitrogen and oxygen respectively.

Ø Trends of Electronegativity : (a) Increases from left to right in a period due to decrease in shared electron nearer to nucleus. (b) Decreases from top to bottom in a group due to shared electron away from nucleus. (c) Highest electronegative atoms in periodic table : F > O > N Ø Effective nuclear change (Zeff) = Nuclear charge – shielding effect Trend : Zeff increases from left to right in a period. Zeff decreases from top to bottom in a group.

qq

Chapter - 4 : Chemical Bonding and Molecular Structure

TOPIC-1General Introduction to Bonding : Ionic and Covalent Bonds

Quick Review Ø Chemical Bond : The attractive force which holds together the constituent particles (atoms, ions or molecules)

in a chemical species is known as chemical bond.

Ø Octet Rule : The octet rule or the electronic theory of chemical bonding was developed by Kossel and Lewis. According to this rule, atoms can combine either by transfer of valence electrons from one atom to another or by sharing their valence electrons. The principle of attaining maximum of eight electrons in the valence shell

of atoms is called octet rule.


Ø Limitations of the Octet Rule : Some limitations of octet rule are : l It cannot be applied to the non-metals after silicon in the periodic table. l These elements can “expand their octet” and have more than eight valence electrons around the central

atom, e.g. PF5, SF6 etc. l Molecules with an odd number of electrons such as NO and NO2 cannot satisty the octet rule. l In some molecules, the central atom cannot possibly have eight valence electrons. For example, LiCl, BeCl2

and BCl3 do not obey the octet rule.

Ø Ionic or Electrovalent Bond : “Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions together is known as ionic bond and electrovalent bond.”

“Electrovalent compounds are those compounds in which constituent particles are ions or positively charged species and they are formed by transfer of electron from a metal to a non-metal.

Ø Properties of Ionic Compounds : Following are the properties of ionic compounds : l Ionic compounds have high melting points and high boiling points. l They have higher enthalpies of fusion and vaporization than molecular compounds. l They are hard and brittle. l These compounds conduct electricity when they are dissolved in water or in molten state. l Ionic solids are good insulators. l The ionic bond is non-directional. This is because each ion is surrounded by oppositely charged ions

uniformly distributed all around the ion.

Ø Covalent Bond : A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. The stable balance of attractive and repulsive forces between atoms when they share electrons is known as covalent bonding.

H + ��

Sharing ofelectrons

OrH H H H—H

The sharing of electrons can also be shown by drawing a circle around each atom. For example, H2 molecule is shown as

Similarly, two chlorine atoms combine with each other to form a molecule of chlorine.

Methane (CH4) :

Depending upon the number of 2, 4 or 6 electrons shared between two atoms, a covalent bond is called single, double or triple bond respectively.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 27 Ø If two atoms share one electron pair, bond is known as single covalent bond and is represented by one dash

(–).

Ø It two atoms share two electron pairs, bond is known as double covalent bond and is represented by two dashes (=).

Ø It two atoms share three electron pairs, bond is known as triple covalent bond and is represented by three dashes (≡).

Ø Lewis Structure of Ionic Compounds :

(i)

(ii)

(iii)

Cl(Mg ) ( )2+

2

–

(iv)

(v)

(vi)

(vii)

(viii)

(ix)

(x)

Ø Lewis Structure of Covalent Compounds :

(i)

(ii)

(iii)

(iv)

(v)

(vi)


(vii)

(viii)

(ix) (x)

(xi) (xii)

(xiii) (xiv)

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 29 Ø Formal Charge : The formal charge of an atom in a polyatomic ion or molecule is defined as – “The difference

between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.

It may be expressed as : Formal charge on an atom in a Lewis structure = (Total number of valence electrons in free atom) – (Total number of non-bonding (lone pair) electrons)

– 12

(Total number of bonding (shared) electrons)

FC = V – L – 12

S

where FC = Formal charge on an atom V = Total number of valence electrons in the free atom L = Total number of electrons present as non-bonding (lone pair) S = Total number of bonding (shared) electrons. Ø Properties of Covalent Compounds : The properties of covalent compounds are : l Low melting and boiling points : Covalent compounds consist of molecules held by weak forces. These

can be easily overcome by heat. l Non-conduction nature : Covalent compounds do not conduct electricity, i.e., electricity does not pass through

the covalent compounds. This is because the covalent compounds do not contain ions or free electrons. l Solubility : Covalent compounds are non-polar and do not dissolve in polar solvents like water. The covalent

compounds however, dissolve in non-polar solvents like benzene, toluene etc. l Slow rate of reaction : The reaction of the covalent compounds are quite slow. The reaction is molecular,

and the molecular reactions are slow. Ø Bond Parameters : The various characteristic features shown by a bond are called the bond parameters. l Bond length – Bond length is defined as “the average distance between the centres of the nuclei of two

bonded atoms in a molecule.”

R = rA + rB

Covalent radius = rA + rB =

d2

[if rA = rB]

Single covalent radius decreases from left to right along a period and increases down a group just like atomic radii.

l Bond angle : It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. For example, the bond angle in some compounds can be represented as under :

30 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI l Bond enthalpy : It is defined as the amount of energy required to break one mole of bonds of a particular

type between two atoms in a gaseous state. The unit of bond enthalpy is kJ mol–1. For example H – H bond enthalpy in hydrogen molecule is 435.8

kJ mol–1. H2(g) → H(g) + H(g) ; ∆H = 435.8 kJ mol–1

l Bond order : In the Lewis description of covalent bond, the bond order is given by the number of bonds between the two atoms in a molecule. Bond order in H2 is one, in O2, it is two and in N2, it is three.

Isoelectronic molecules and ions have identical bond orders. For example F2 and O 22– have bond order

1 and N2, CO and NO+ have bond order 3. The stabilities of molecules can be understood by the statement ”with increase in bond order, bond enthalpy

increases and bond length decreases.” Ø Types of Covalent Bond : Depending upon the nature of combining atoms, the covalent bonds can be of two

types : l Non-polar Covalent Bond : This type of bond is formed between the two atoms of the same element.

Atoms of the same element attract electron equally. So, in other words, the shared electron pair will lie exactly midway between the two atoms. This type of covalent bond is described as a non-polar covalent bond.

For example : H2, O2 and Cl2 etc. l Polar covalent Bond : This type of bond is formed between two atoms of different elements; the shared

pair of electrons does not lie exactly midway between the two atoms. In fact, it lies more towards the atom which is more electronegative. The atom with higher affinity for electrons develops a slight negative charge and atom with lesser affinity for electrons develop a slight positive charge. For example : HCl, H2O etc.

Ø Dipole Moment : It is defined as “the product of the magnitude of the charge and the distance of separation between the charges.”

Dipole moment (µ) = Charge (q) × Distance of separation (d) It is usually expressed in Debye (D). Dipole moment is also expressed in the units of C m : 1D = 3.33564 × 10–30 C m The dipole of HCl may be represented as :

H Cl

Know the Terms Ø Valence Electrons : The outermost shell electrons which take part in the formation of chemical bond are called

valence electrons. Ø Lewis Symbols : G.N. Lewis introduced simple notations to represent

valence electrons in an atom, which are called Lewis symbols, e.g. Li·

, Be·

·.

Ø Octet Rule : The rule by which atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to complete an octet in their valence shell is called octet rule.

Ø Ionic Bond or Electrovalent Bond : The bond formed, as a result of the electrostatic attraction between the positive and negative ions was termed as the ionic bond or electrovalent bond.

Ø Covalent Bond : The bond formed as a result of sharing of electron pairs between atoms is called covalent bond. Ø Single Covalent Bond : When two atoms share one electron pair, they are said to be joined by a single covalent

bond. Ø Double Bond : If two atoms share two pairs of electrons, the covalent bond between them is called a double bond. Ø Triple Bond : If two atoms share three pairs of electrons, the covalent bond between them is called a triple bond. Ø Lattice Enthalpy : The lattice enthalpy of an ionic solid is defined as the energy required to completely separate

one mole of a solid ionic compound into gaseous constituent ions. Ø Bond Length : It is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. Ø Bond Angle : It is defined as the angle between the orbitals containing bonding electron pairs around the central

atom in a molecule/complex ion. Ø Bond Enthalpy : It is defined as the amount of energy required to break one mole of bonds of a particular type

between two atoms in a gaseous state. Ø Electrovalency : The valence of an ion, equal to the number of positive or negative charges acquired by an atom

through a loss or gain of electrons is known as electrovalency.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 31 Ø Polarizability : The power of an ion to distort the other ion is polarizing power and the tendency of an ion to get

distorted is called polarizability. Ø Covalency : The bond formed between two atoms by mutual sharing of electrons between them so as to complete

their octets is known as covalent bond and number of electrons involved is covalency.

Know the Facts Ø Under normal conditions, no other element exists as an independent atom in nature, except noble gases. Ø Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. Ø In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated

by the noble gases. Ø The formation of a negative ion from a halogen atom and a positive ion from an alkali metal atom is associated

with the gain and loss of an electron by the respective atoms. Ø The negative and positive ions are stabilized by electrostatic attraction. Ø The Lewis dot structures provide a picture of bonding in molecules and ions in terms of the shared pair of

electrons and the octet rule. Ø Generally the lowest energy structure is the one with the smallest formal charges on the atoms. Ø The formal charge is a factor based on a pure covalent view of bonding in which electron pairs are shared equally

by neighbouring atoms.

TOPIC-2Resonance and VSEPR Theory

Quick Review Ø Resonance : When a molecule cannot be represented by a single structure but its characteristic properties can

be described by two or more than two structures, then actual structure is said to be a resonance hybrid of these structure.

Ø Conditions for Writing Resonating Structure : The essential conditions for writing resonating structures are given as below :

l The contributing structures should have same atomic positions. l These structures should have same number of unpaired electrons. l These structures should have nearly same energy. l The structures should be so written that negative charge is present on an electronegative atom and positive

charge is present on an electropositive atom. l In these structures, the like charges should not reside on adjacent atoms.

Ø Resonance stabilizes the molecule because the energy of the resonance hybrid is less than the energy of any single canonical structures.

Ø Resonance Energy : It is the difference between the actual bond energy of the molecule and that of the most stable resonating structures (having least energy). Thus,

Resonance energy = Actual bond energy – Energy of the most stable resonating structure

Ø Valence Shell Electron Pair Repulsion (VSEPR) Theory : In 1940, Sidgwick and Powell proposed a simple theory based on the repulsive interactions of the electron pairs in the valence shell of the atoms. It was further developed and refined by Nyholm and Gillespie in 1957.

They suggested that the shapes of molecules can be determined by the number of electron pairs (bonding as well as non-bonding) in the valence shell of the central atom.

This suggestion is the basis of valence shell electron pair repulsion (VSEPR) theory.

The basic idea of this theory is that bonded atoms in a molecule adopt that particular arrangement in space around the central atom which keeps them on the average as apart as possible.

Ø Postulates of VSEPR Theory : l The geometry of a molecule depends upon the total number of valence shell electron pairs (bonded or

non-bonded) around the central atom in the molecule.

32 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI l Repulsion between valence shell electron pairs determines molecular shape. In order to minimize electron-

electron repulsion or to acquire a state of minimum energy or maximum stability, the valence shell electron pairs keep themselves as far apart as possible.

l In case more than one geometrical arrangements are possible for a given number of bonded and non-bonded pairs, then the most stable arrangement are consistent with the following hypothesis :

(a) Repulsion between a lone pair and lone pair of electrons is different from that between the bond pairs or one lone pair and one bond pair. The repulsive interactions decrease in the order : Lone pair-Lone pair > Lone pair-Bond pair > Bond pair-Bond pair.

The presence of lone pairs in addition to bond pairs will result in certain distortions in the regular geometry of molecules.

(b) Repulsive forces decrease sharply with increasing angle between the electron pairs. They are strong at 90°, much weaker at 120° and very weak at 180°.

Thus, the molecules in which the central atom is surrounded only by similar bonded electron pairs will have regular geometries while those in which the central atom is surrounded by bond pairs as well as lone pairs will have irregular geometries.

Effect of lone pair of electrons (a) Four orbitals having bond pair electrons (b) three orbitals with bond pair electron and one orbital with lone pair electrons (angle between orbitals with bond pair electrons decreases)

Know the Terms Ø Resonance : The delocalization of a pair of electrons in a molecule or ion which results in observed bond length,

bond order, bond energy different from normal covalent bond is called resonance. Ø Resonance Hybrid : The structure which is near to all resonating structures and nearly explain the property of

that molecule or ion is called resonance hybrid.

Know the Facts Ø The shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded)

around the central atom. Ø The repulsive interaction of electron pairs decrease in the order – lp – lp > lp – bp > bp – bp (where, lp = lone pair, bp = bond pair)

TOPIC-3Valence Bond Theory, Hybridisation, Molecular Orbital Theory and Hydrogen Bonding

Quick Review Ø According to V.B.T., a covalent bond is formed between the two atoms by the overlap of half filled valence atomic

orbitals of each atom containing one unpaired electron. Ø The bond is formed when two atoms approach each other in such a way that occupied orbitals with similar

energies are able to overlap. Ø Greater the overlap, stronger is the bond. Ø Types of Overlapping and Nature of Covalent Bonds : Depending upon the type of overlapping, the covalent bonds may be divided into two types : (a) Sigma (s) bond and (b) Pi (p) bond (a) Sigma (s) bond : This type of covalent bond is formed by the end to end (head on) overlapping of bonding

orbitals along the internuclear axis. The overlap is known as head on or axial overlap. The sigma bond is formed by any one of the following types of combinations of atomic orbitals :

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 33 (i) s-s overlapping : In this type, two half-filled s-orbitals overlap along the internuclear axis as shown below.

(ii) s-p overlapping : This type of overlapping between the half-filled s-orbital of one atom and p-orbital of the other atom.

s-orbital p-orbital s-p ovrelap

(iii) p-p overlapping : This type of overlapping occurs between half-filled p-orbitals of the two approaching atoms.

p-orbital p-orbital p-p-overlap

(b) Pi (p) bond : This type of covalent bond is formed by the sidewise overlap of the half-filled atomic orbitals of bonding atoms. Such an overlap is known as sidewise or lateral overlap. The atomic orbitals overlap in such a way that their axis remain parallel to each other and perpendicular to the internuclear axis. The orbital obtained as a result of sidewise overlap consists of two saucer type charged clouds above and below the plane of the participating atoms.

or

�

�

p-orbital p-orbital p-p-overlap

Ø Strength of sigma and pi-bonds : The strength of a covalent bond depends upon the extent of overlapping of atomic orbitals forming the bond. During the formation of a sigma bond, the overlapping of orbitals takes place to a larger extent.

Ø On the other hand, during the formation of a pi-bond, the overlapping occur to a smaller extent. Therefore, a sigma bond is strong than a pi-bond.

Ø Hybridisation : According to Pauling, to account for the shapes of polyatomic molecules, the atomic orbitals combine to form new set of equivalent orbitals known as hybrid orbitals, are used in bond formation. This process is known as hybridisation. This may be defined as—

“The process of intermixing of the orbitals of slightly different energies so as to redistribute their energies resulting in the formation of new set of orbitals of equivalent energies and shape.”

Ø Characteristics of Hybridisations : l The number of hybridised orbitals is equal to the number of the orbitals that get hybridised. l The hybridised orbitals are always equivalent in energy and shape. l The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals. l The hybrid orbitals are directed in space in some preferred directions to have stable arrangement. Therefore,

the type of hybridisation indicates the geometry of the molecule. Ø Conditions for Hybridisations : l Only the orbitals present in the valence shell of the atom are hybridised. l The orbitals undergoing hybridisation should have only a small difference in energy. The orbitals which

differ largely in energy cannot take part in hybridisation. l Promotion of electron is not essential condition prior to hybridisation. l It is not essential that only half filled orbitals participate in hybridisation. In certain cases, even filled orbitals

of valence shell participate in hybridisation.

34 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI There are many types of hybridisation involving different atomic orbitals. The different types of hybridisation are

discussed below : Ø Types of Hybridisation : These are given as below : l sp Hybridisation : In sp hybridisation, one s and one p orbital mix and form two equivalent sp hybrid orbitals.

The orbitals suitable for sp hybridisation are s and pz orbitals, if the hybrid orbitals are to lie along Z-axis. Each sp hybrid orbitals has 50% s-character and 50% p-character. These two sp hybrid orbitals are oriented in

opposite direction forming an angle of 180°. The sp-hybridisation is pictorially shown as below :

Formation of sp hybrid orbitals

Example : BeF2, BeCl2 etc.

l sp2 Hybridisation : The combination of one s and two p orbitals gives rise to three sp2 hybrid orbitals. These

three hybrid orbitals lie in the same plane and are directed towards the three corners of an equilateral triangle

in a plane having an angle of 120°. Each sp2 hybrid orbital has 33·3% s-character and 66·7% p-character.

s

pzpy

px

sp2-hybridisation

px -orbital

px

Three -hybridorbitals

sp2

120°

Formation of sp2 hybrid orbitals Example : BCl3, BF3 etc.

l sp3 Hybridisation : In sp3 hybridisation, one s and three p orbitals mix and form four sp3 hybrid orbitals of equivalent energies and shape. Each sp3 hybrid orbitals has 25% s-character and 75% p-character. The four sp3 hybrid orbitals are directed towards the four corners of the tetrahedron. The angle between sp3 hybrid orbital is 109°28’.

z

yx

One -orbitals Three -orbitalsp

x

z

y

Fourhybrid orbitals

sp3

109°28'sp3-hybridisation

Formation of sp3 hybrid orbitals

Example : CH4, CCl4 etc.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 35 l sp3d Hybridisation : This hybridisation involves the mixing of one s, three p and one d-orbitals. These five

orbitals hybridise to form five sp3d-hybrid orbitals. The mixing of five orbitals is shown in figure. These hybrid orbitals point towards the corners of a trigonal bipyramidal geometry.

3s

3px 3py

y

x

z

3pz

3dz2

x

sp3d hyrbid orbitals havingtrigonal bipyramidal

geometry

Formation of sp3d hybrid orbitals Example : Phosphorus pentachloride, PCl5. l sp3d2 Hybridisation : A combination of one s, three p and two d atomic orbitals (dx2–y2, dz2) leads to six sp3d2

hybrid orbitals which are directed octahedrally with bond angle of 90°. Example: In SF6, one electron each from 3s and 3px orbitals is promoted to 3d-orbitals as shown in Fig. (b).

These six orbitals get hybridised to form six sp3d2 hybrid orbitals. Each of these sp3d2 hybrid orbitals overlaps with 2p-orbitals of fluorine to form S—F bond. Thus, SF6 molecule has octahedral structure as shown in Fig. (b).

(a) S (Ground state)

(b) S (Excited state)

(c) SF6

3s

sp3 2d hyrbidisation

3p 3d

F F F F F F

S

F

F

F

FF

F

Octahedral geometry

(a) (b)

(a) Formation of SF6 molecule involving sp d -hybridisation (b) Octahedral geometry of SF molecule3 2

6

l sp3d3 Hybridisation : This involves the mixing of one s, three p and three d-orbitals forming seven sp3d3-hybrid orbitals having pentagonal bipyramidal geometry. The geometry of IF7 molecule can be explained on the basis of sp3d3 hybridisation.

Ø Geometry of IF7 molecule : The outer electronic configuration of iodine atom is 5s25p5. To make seven bond, one electron each from s and p-orbitals are promoted to the higher vacant 5d-orbitals as shown in Fig. (a) and (b). These seven orbitals are then hybridised to give seven sp3d3-hybrid orbitals. Each of these sp3d3 hybrid orbitals overlaps with 2p-orbitals of fluorine to form IF7 molecule having pentagonal bipyramidal geometry.

(a) I (Ground state)

(b) I (Excited state)

(c) IF7

3s

sp3 3d hyrbidisation

3p 3d

F F F F F F

I

F

F

F

F

F

F

Pentagonal bipyramidalF

(b)(a)

F

(a) Formation of IF7 molecule involving sp3d3-hybridisation (b) Pentagonal bipyramidal geometry of IF7 molecule

36 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI In this geometry, all the bond angles are not equal. Five F-atoms are directed towards the vertices of a regular

pentagon making an angles of 72°. The other two F-atoms are at right angles (90°) to the plane. Due to different bond angles, the bonds are different in length. The axial bonds are larger than equatorial bonds.

Ø Molecular Orbital Theory : Molecular orbital theory was proposed by F. Hund and R.S. Mulliken in 1932. The basic idea of molecular orbital

theory is that atomic orbitals of individual atoms combine to form molecular orbitals. The electrons in molecules are present in the molecular orbitals which are associated with several nuclei. These

molecular orbitals are filled in the same way as the atomic orbitals in atoms are filled. Salient Features of Molecular Orbital Theory : l Just like an atom, molecule has orbitals of definite energy levels. Like electrons of atoms are present in atomic

orbitals, electrons of a molecule are present in different molecular orbitals. l Molecular orbitals are formed by the combination of atomic orbitals of proper symmetry and comparable

energies. The atomic orbitals of these atoms merge into molecular orbitals. l An electron in an atomic orbital is under the influence of only one nucleus. However, an electron in a

molecular orbital is under the influence of two or more nuclei depending upon the number of atoms present in the molecule.

l The number of molecular orbitals formed is equal to the number of combining atomic orbitals. When two atomic orbitals combine, they form two molecular orbitals. These are called bonding molecular orbital and antibonding molecular orbital.

l The bonding molecular orbital has lower energy and hence greater stability whereas corresponding antibonding molecular orbital has more energy and hence lesser stabilty.

l The molecular orbital formed by the addition of two atomic orbitals is called bonding molecular orbital. It is shown in following figure :

++

+

1s 1s

+ + +

+

�1sBonding molecular orbitalAtomic orbital

Formation of bonding molecular orbital

l The molecular orbital formed by the subtraction of two orbitals is called antibonding molecular orbital. It is shown in following figure :

+–

+

1s 1s

+ –

��1sAntibonding molecular orbital

Atomic orbital Node

+ –

Formation of antibonding molecular orbital

It is clear from the figure that the electron density in between the two nuclei is practically zero and is concentrated in regions away from each nucleus. In this case, the electron probability densities of two atomic orbitals get cancelled in the centre (by subtraction) so that there is no probability of finding the electron in the region of overlap i.e. between the nuclei. This situation does not favour the bond formation. Such an orbital is called antibonding molecular orbital. It is designated as s*1s (called sigma star 1s). The asterisk ‘*‘ is used to represent antibonding molecular orbital.

Ø Conditions for the Combination of Atomic Orbitals : Molecular orbitals are formed by the combination of atomic orbitals. But all types of orbitals cannot combine with each other. The following are the main conditions for effective combination of atomic orbitals.

l The combining atomic orbitals must have same or nearly the same energies. l The extent of overlapping between the atomic orbitals of two atoms should be large.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 37 Ø Types of Molecular Orbitals : l Bonding molecular orbitals : A molecular orbital that is obtained by the additon overlap (i.e, when the lobes

of atomic orbital overlap with the same sign) of two atomic orbitals is known as bonding molecular orbital. The energy of bonding molecular orbital is less than that of the constituent overlapping atomic orbitals. The difference in energy between the combining atomic orbitals and the bonding molecular orbital formed, is called the stabilisation energy, because, it stabilizes the molecule.

l Non-bonding molecular orbitals : Orbital, which are completely inside the closed shell of atom, do not play any part in the formation of molecule. They remain unaffected and are termed non-bonding molecular orbitals. When atomic orbitals do not interact with each other, formation of non-bonding molecular orbitals takes place. The reasons for no interaction may be one of the following :

(i) One orbital is far away from other orbital, so overlapping is not possible. (ii) Symmetries of atomic orbitals are not same. (iii) Energies of atomic orbitals are not same. l Antibonding molecular orbitals : A molecular orbital that is obtained by the subtraction overlap (i.e., when

the lobes of atomic orbitals overlap with the opposite sign) of two atomic orbitals is known as antibonding molecular orbital. The energy of antibonding molecular orbital is higher than that of the constituent overlapping atomic orbitals. The difference in energy between the antibonding molecular orbital and the combining atomic orbital is known as destabilisation energy, because it destabilizes the molecule.

For molecules O2 onwards, the increasing order of energies for molecular orbitals in which they are filled is as follows :

s1s, s*1s, s2s, s*2s, s2pz, (p2pz = p2py), (p*2pz = p*2py), s*2pz Ø However, for homonuclear diatomic molecules of second row elements such as, Li2, Be2, B2, C2, N2 the s2pz MO

is higher in energy than p2px and p2py MOs. The increasing order of energies order for MOs for their molecules is s1s, s*1s, s2s, s*2s, (p2px = p2py), s2pz, (p*2px = p*2py), s*2pz Ø The filling of molecular orbitals take place by following the rules of Aufbau principle, Pauli’s exclusion principle

and Hund’s rule in the similar way as for atomic orbitals. Ø Presence of unpaired electron shows that molecule is paramagnetic, however, paired electrons show that it is

diamagnetic. Ø Hydrogen Bonding : When hydrogen atom is bonded to atoms of highly electronegative elements such as

fluorine, oxygen or nitrogen, the hydrogen atom forms a weak bond with the electronegative atom of the other molecule. This weak bond is called hydrogen bond. For example, in hydrogen fluoride, HF, hydrogen atom forms a weak bond with fluorine atom of the high bonding molecule while remaining bonded to its fluorine atom. This may be shown as :

H—F ................ H—F ................. H—F ................ In other words, hydrogen atom acts as a bridge between two atoms, holding one atom by a covalent bond and

the other atom by a hydrogen bond. The hydrogen bond is represented by dotted line (...........) while the covalent bond is represented by solid line (––).

Ø Types of Hydrogen Bond— Hydrogen bonds can be classified into two types : (a) Intermolecular hydrogen bond (b) Intramolecular hydrogen bond (a) Intermolecular hydrogen bond : Intermolecular hydrogen bond is formed between different molecules of the

same or different substances. For example, (i) hydrogen bond between the molecules of hydrogen fluoride. (ii) hydrogen bond in alcohol or water molecules. (b) Intramolecular hydrogen bond : Intramolecular hydrogen bond is formed between the hydrogen atom and

the highly electronegative atom (F, O or N) present in the same molecule. Intramolecular hydrogen bond results in the cyclisation of the molecule and prevents its association. Consequently,

the effect of intramolecular hydrogen bond on the physical properties is negligible.

38 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Applications of Intermolecular H-bonding : l Melting point and boiling point of water : Water has the lowest molecular weight among the hydrides of

group 16 elements yet it has the highest melting and boling points. It is due to intermolecular H–bonding in H2O.

l Ice has less density than water : In crystal structure of ice, every water molecule is associated with four other water molecules by H-bonding in a cage like tetrahedral structuure. On melting, H-bonds in the ice are broken and space between water molecules decreases and density of water increases up to 4°C. Above 4°C, more H-bond are broken. The water molecules get apart from each other and the density again decreases. Thus, water has maximum density at 4°C.

Ø Applications of Intramolecular H-Bonding : l Volatile character of nitrophenols : ortho-nitrophenol is more volatile (b.p. 214°C) as compared to meta (b.p.

290°C) and para (b.p. 279°C). It is due to chelation.O

N

H

O

O

��

��

l In meta and para isomer, chelation is not possible due to the formation of desired size of ring.

Know the Terms Ø Sigma (s) bond : The covalent bond which is formed by the end to end (head on) overlap of bonding orbitals

along the internuclear axis is called sigma (s) bond. Ø Pi (p) bond : The covalent bond which is formed by the sidewise overlap of bonding orbitals perpendicular to the

internuclear axis is called pi (p) bond. Ø Hybridisation : It is defined as the process of intermixing of the orbitals of slightly different energies so as to

redistribute their energies resulting in the formation of new set of orbitals of equivalent energies and shape. Ø Bonding Molecular Orbital : The molecular orbital which is formed by addition of two atomic orbitals is called

bonding molecular orbital. Ø Anti-bonding Molecular Orbital : The molecular orbital which is formed by subtraction of two atomic orbitals is

called anti-bonding molecular orbital. Ø Bond Order : It is defined as one half the difference between the number of electrons present in the bonding and

the anti bonding orbitals. [Bond order = 12

(Nb – Na)]

Ø Hydrogen Bond : It can be defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) of another molecule.

Know the Facts Ø The extent of overlap decides the strength of a covalent bond. Generally, greater the overlap, the stronger is the

bond formed between two atoms. Ø According to orbital overlap concept, the formation of a covalent bond between two atoms results by pairing of

electrons present in the valence shell having opposite spins. Ø When orbitals of two atoms come close to form bond, their overlap may be positive, negative or zero depending

upon the sign (phase) and direction of orientation of amplitude of orbital wave function in space. Ø According to concept of hybridization, the atomic orbitals combine to form new set of equivalent orbitals known

as hybrid orbitals. Ø The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised. Ø The hybridised orbitals are always equivalent in energy and shape. Ø The orbitals present in the valence shell of the atom are hybridised. Ø The orbitals undergoing hybridization should have almost equal energy. Ø The sigma (s) molecular orbitals are symmetrical around the bond axis whereas pi (p) molecular orbitals are not

symmetrical. Ø The increasing order of energies of various molecular orbitals in case of O2 or F2 : s1s < s*1s < s2s < s*2s < s2pz < (p2px = p2py) < (p*2px = p*2py) < s*2pz

ØThe increasing order of energies of various molecular orbitals in case of Li2, Be2, B2, C2 or N2 :

s1s < s*1s < s2s < s*2s < (p2px = p2py) < s2pz < (p*2px = p*2py) < s*2pz

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 39 Ø The distribution of electrons among various molecular orbitals is called the electronic configuration of the

molecule. Ø The bond length decreases as bond order increases. Ø If all the molecular orbitals in a molecule are doubly occupied, the substance is said to be diamagnetic (repelled

by magnetic field). Ø If one or more molecular orbitals are singly occupied, the substance is said to be paramagnetic (attracted by

magnetic field). Ø The bond strength increases as bond order increases. Ø The bond dissociation energy increases as bond order increases. Ø Density of H2O is maximum at 4° due to strongest intermolecular hydrogen bonding. Ø Bond order zero means no possibility of that molecule.

Ø Bond order = 12

(Nb – Na)

Where Nb = No. of electrons in bonding molecular orbital Na = No of electrons in anti-bonding molecular orbital.

qq

Chapter - 5 : States of Matter Gases, Liquids and Solids

TOPIC-1Factors Governing States of Matter and Gas Laws

Quick Review Ø Matter is classified into three states : Solids, Liquids and Gases which can be interchanged by changing

temperature. Ø Solids have definite mass, volume and shape, high density and are incompressible. Ø Liquids have definite mass and volume but not definite shape. Their densities are less than solid and have very

little compressibility. Ø Gases have definite mass but not finite volume and shape and are highly compressible and low density. Ø Intermolecular forces are the strongest in solids and weakest in gases. Ø Intermolecular distances are the shortest in solids and the longest in gases. Liquids show intermediate behaviour. Ø Intermolecular forces which hold the molecules together may be London or Dispersion forces, dipole-dipole

forces, dipole-induced dipole forces and hydrogen bonding. Ø The states of matter depend upon their intermolecular forces and thermal energy. Intermolecular forces bind

the molecules together, however thermal energy keep them apart. Therefore, at sufficient low temperature, the attractive intermolecular forces dominate and solid state is the stable state. However, at high temperatures, thermal forces dominate and gaseous state is the stable state. Liquid state exists at intermediate temperatures.

Ø Of the three physical states of matter, the gaseous state is most random state and solid state is least random.Ø Gas Laws : The generalisations regarding the behaviour of gases.Ø Boyle’s Law : According to this law, ”The volume of a fixed mass of a gas is inversely proportional to the pressure

at constant temperature.” Boyle’s Law Equation : (At constant T)

or PV

∝1

or P1V1 = P2V2

Ø Charle’s Law : According to this law, ”The volume of fixed mass of a gas is directly proportional to the temperature (in Kelvin) at constant pressure.” The volume of a given amount of a gas at constant pressure increases (or decreases) by a constant fraction (1/273.15) of its volume at 0°C for each degree rise (or fall) in temperature.

Charle’s Law Equation : (At constant P, n)

V ∝ T or V1/T1 = V2/T2

Ø Gay Lussac’s Law : According to this law, “The pressure of fixed mass of a gas at constant volume is directly proportional to its temperature in Kelvin.”

Gay Lussac’s Law Equation : (At constant V, n) P ∝ T

1

1

PT

=

2

2

PT

40 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XIØ Avogadro’s Law : According to this law, “Under similar conditions equal volumes of various gases contain equal

number of molecules.” Avogadro’s Law Equation : (At constant T, P) V ∝ n V = k4n (where n = m/M)Ø Dalton’s Law of Partial Pressures : According to this law, “The pressure of the mixture of non-reacting gases is

equal to the sum of their partial pressure when enclosed in the same volume under similar conditions.” Dalton’s Law Equation : P = P1 + P2 + P3Ø Graham’s Law of Diffusion : This law states that, “The rate of diffusion of gases (V/t) is inversely proportional to

square root of their densitities under similar conditions.”

Graham’s Law : r ∝

1d

1

2

rr

=

2 2

1 1

MM

dd

=

Ø Ideal gas equation, PV = nRT gives the relationship between different parameters and is obtained by combining Boyle’s law and Charle’s law. Ideal Gas Equation : PV = nRT

1 1

1

P VT

=

2 2

2

P VT

Ø Absolute Zero : The hypothetical temperature at which all the gases are supposed to occupy zero volume. Its numerical value on Celsius scale is – 273.15 °C.

Ø Vapour Density of a Gas : It is the ratio of the mass of certain volume of gas to the mass of same volume of hydrogen under similar conditions.

Molecular mass of gas = 2 × Vapour density

Know the Terms Ø Intermolecular Forces : The forces of attraction and repulsion between interacting particles (atoms and molecules)

are called intermolecular forces. Ø Van der Waals Forces : Attractive intermolecular forces are known as van der Waals forces, in honour of Dutch

scientist Johannes van der Waals, who explained the deviation of real gases from the ideal behavior through these forces.

Ø London Forces : The force of attraction between two temporary dipoles i.e. between two non-polar molecules is called London force. It is also known as dispersion force.

Ø Dipole-dipole Forces : These forces act between the molecules having permanent dipole i.e. between polar molecules.

Ø Dipole-Induced dipole Forces : These forces operate between the polar molecules having permanent dipole and the molecules lacking permanent dipole i.e. between a polar and a non-polar molecule.

Ø Thermal Energy : It is the energy of a body arising from motion of its atoms or molecules. Ø Absolute Temperature Scale : It is defined as a new scale of temperature such that t°C on new scale is given by

T = 273.15 + t and 0°C will be given by To = 273.15. It is also called Kelvin scale temperature or thermodynamic scale temperature.

Ø Absolute Zero : The lowest hypothetical or imaginary temperature at which gases are supposed to occupy zero volume is called absolute zero.

Ø Avogadro Constant : The number of molecules in one mole of a gas has been determined to be 6.022 × 1023 and is called Avogadro constant.

Ø Ideal Gas : A gas that follows Boyle’s law, Charle’s law and Avogadro law strictly is called an ideal gas. Ø Partial Pressure : In a mixture of gases, the pressure exerted by the individual gas is called partial pressure. Ø Aqueous Tension : Pressure exerted by saturated water vapour is known as aqueous tension. Ø STP (Standard Temperature and Pressure) : STP means 273.15K (0°C) temperature and 1 bar (i.e. exactly 105

pascal pressure. At STP, molar volume of an ideal gas or combination of ideal gases is 22.7 L mol–1.

Know the Facts Ø Chemical properties of a substance do not change with the change of its physical state, but rate of chemical

reactions do depend upon the physical state. Ø Intermolecular forces are the forces of attraction and repulsion between interacting particles (atoms and

molecules). Ø Attractive intermolecular forces are known as van der Waals forces.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 41 Ø Induced dipole moment depends upon the dipole moment present in the permanent dipole and the polarisability

of the electrically neutral molecule. Ø High polarisability increases the strength of attractive interactions. Ø Energy of hydrogen bond varies between 10 to 100 kJ mol–1. Ø Strength of the hydrogen bond is determined by the columbic interaction between the lone pair of electrons of

the electronegative atom of one molecule and the hydrogen atom of other molecule. Ø Thermal energy is a measure of average kinetic energy of the particles of the matter and is responsible for

movement of particles. This movement of particles is called thermal motion. Ø When thermal energy of molecules is reduced by lowering of temperature, the gases can be very easily liquefied. Ø Predominance of thermal energy and the molecular interaction energy of a substance in three states is represented

as follows :

Gas → Liquid → Solid

Predominance of intermolecular interaction.

Gas ← Liquid ← Solid

Predominance of thermal energy

Ø Gases are highly compressible. Ø Gases exert pressure equally in all directions. Ø Standard Ambient Temperature and Pressure (SATP) : It means 298.15 K temperature and 1 bar (i.e. exactly 105

Pa). At SATP, the molar volume of an ideal gas is 24.789 L mol–1. Ø Ideal gas equation is a combination of Boyle’s law, Charle’s law and Avogadro law. Ø PDry gas = PTotal – Aqueous tension

TOPIC-2Kinetic Theory of Gases and Behaviour of Real Gases

Quick Review Ø Kinetic molecular theory explains the behaviour of ideal gas. Ø Under ordinary conditions, real gases show deviations from the ideal gas behaviour. Real gases follow ideal gas

behaviour at high temperatures and very low pressures. Ø The two main causes of these deviations are the finite size of gas molecules and the existence of attractive

intermolecular forces between them. Taking into account these facts, van der Waals modified the ideal gas equation and gave the modified equation known as van der Waals equation.

( ) =2

2P+ (V – ) RTVn a nb n

where

2Va

is the pressure correction term and b, the volume correction term takes account the finite volume of

the gas molecules. Ø The deviation from ideal behaviour can be measured in terms of compressibility factor Z,

Z =

PVRTn

or Z =

real

ideal

VV

Ø All the gases can be liquefied by compressing them and cooling them. Ø Critical temperature is the temperature above which it cannot be liquefied even at high pressure. Ø The minimum pressure required to liquefy a gas at its critical temperature is known as critical pressure. Ø The volume occupied by one mole of a gas at its critical temperature and under its critical pressure is known as its

critical volume. Ø The critical temperature, pressure and volume are called critical constants.

a = 2 8

or =RT RT

vM Mp

u = 3 3RTM

PVnM

or

42 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø K.E.of gas ∝ T

K.E. of N molecules of the gas = 32RTNNA

Know the Terms Ø Compressibility Factor : It is the ratio of the product PV and nRT. Ø Boyle’s Temperature or Boyle’s Point : The temperature at which a real gas obeys ideal gas law over an appreciable

range of pressure is called Boyle’s temperature or Boyle’s point. Ø Critical Temperature : It is the temperature of a gas in its critical state, above which it cannot be liquefied by

pressure alone. Ø Critical Pressure : It is the pressure of a gas or vapour in its critical state. The critical pressure of a substance is the

pressure required to liquefy a gas at its critical temperature.

Know the Facts Ø Gases consist of large number of identical particles (atoms or molecules) that are so small and so far apart on the

average that the actual volume of the molecules is negligible in comparison to the empty space between them. They are considered as point masses. This assumption explains the great compressibility of gases.

Ø There is no force of attraction between the particles of a gas at ordinary temperature and pressure. It explains that gases expand and occupy all the space available to them.

Ø Particles of a gas are always in constant and random motion. So, a gas has no fixed shape. Ø Particles of a gas move in all possible directions in straight lines. During their random motion, they collide with

each other and with walls of the container. Ø Pressure is exerted by the gas as a result of collision of the particles with the walls of the container. Ø Collision of gas molecules are perfectly elastic. Ø At any particular time, different particles in the gas have different speeds and hence different kinetic energy. Ø Real gases do not follow Boyle’s law, Charle’s law and Avogadro law perfectly under all conditions. Ø Real gases show deviations from ideal gas law because molecules interact with each other. Ø The pressure exerted by the gas is lower than the pressure exerted by the ideal gas.

Pideal = Pan

real(ObservedPressure)

(Correction term)

+2

2V

Ø Real gases show ideal behavior when conditions of temperature and pressure are such that the intermolecular forces are practically negligible.

Ø The real gases show ideal behavior when pressure approaches zero. Ø For ideal gas, the value of compressibility factor, Z = 1 at all temperatures and pressures because PV = nRT. Ø For gases which deviate from ideality, value of Z deviates from unity. Ø Atverylowpressure,Z≈1forallgasesandbehaveasidealgas. Ø At high pressure, Z > 1 for all gases, which are more difficult to compress. Ø At intermediate pressure, Z < 1 for most gases. Ø A gas below its critical temperature can be liquefied by just applying pressure and is called the vapour of the

substance.

TOPIC-3Physical Properties of Liquid States

Quick Review Ø Liquids may be considered as continuation of gas phase into a region of small volume and very strong molecular

attractions. Ø Some properties of liquids, e.g., surface tension and viscosity are due to strong intermolecular attractive forces. Ø Vapour Pressure—The pressure exerted by the vapours of the liquid in equilibrium with the liquid at particular

temperature. Vapour pressure of a liquid depends upon temperature. It increases with rise in temperature. Ø Boiling Point of Liquid—The temperature at which, the vapour pressure of the liquid becomes equal to the

atmospheric pressure (or pressure applied on the surface of liquid) is called boiling point of the liquid.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 43 Ø If pressure is 1 atmosphere, the boiling point is called the normal boiling point of the liquid. Ø If pressure is 1 bar, the boiling point is called the standard boiling point of the liquid. Ø The value of standard boiling point is slightly lower than that of normal boiling point. Ø The amount of heat required to change one mole of the liquid into its vapours at the boiling point is known as

heat of vaporization of the liquid. Ø Surface tension: It is defined as the force acting per unit length perpendicular to the line drawn on the surface

of liquid. It is responsible for many characteristic properties of liquids like curved meniscus, rise in a capillary, spherical shape of liquid drops.

Unit : kgs–2 or Nm–1

Surface tension = Force ( )

Length ( )Fl

Ø Liquids tend to minimize their surface area due to surface tension. Ø The energy required to increase the surface area of the liquid by one unit is called surface energy. Its unit is Jm–2. Ø Liquids tend to rise or fall in the capillary due to surface tension. The height of liquids in the capillary is given by.

h = 2Tr l gcosq

Where T = Surface tension q = Contact angle r = Radius of capillary l = Density of liquid Ø Viscosity : It is the internal force of friction of a liquid. It depends upon the strength of intermolecular forces and

shape of molecules. Unit : Ns–2 or Pa S or kg m–1s–1 or poise Ø Surface tension and viscosity both decreases with rise in temperature. Ø When a liquid flows over a fixed surface, the layer of molecules in the immediate contact of surface is stationary.

The velocity of upper layers increases as the distance of layers from the fixed layer increases. This type of flow in which there is a regular gradation of velocity in passing from one layer to the next is called laminar flow. If we select any layer in the flowing liquid, the layer above it accelerates its flow and the layer below this retards its flow.

Gradetion of velocity in the laminar flow

Ø If the velocity of the layer at a distance dz is changed by a value du then velocity gradient is given by the amount dudz

.

Ø A force is required to maintain the flow of layers. This force is proportional to the area of contact of layers and velocity gradient.

i.e. F ∝ A (A = area of contact)

Fdudz

∝ (dudz

Velocity gradient i.e. change in velocity with distance)

F Adudz

∝

∴ F Adudz

= η

Where h = Proportionality constant (coefficient of viscosity)

Ø S.I. unit of viscosity = NSm–2

Or Pascal second Or Kg m–1 s–1

44 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Factors affecting Viscosity :

(i) Temperature : Viscosity1

temperaturea

(ii) Pressure : Viscosity a pressure (iii) Hydrogen bonding : Viscosity a extent of hydrogen bonding

Know the Terms Ø Liquid State : It is the state of matter in which the intermolecular forces of attraction are strong enough to keep

the molecules together. Ø Vapour Pressure or Equilibrium Vapour Pressure : It is the pressure exerted by the vapours in equilibrium with

liquid at a fixed temperature. Ø Boiling Point : It is the temperature at which vapour pressure of liquid becomes equal to the external pressure of

liquid. Ø Surface Tension : It is the force acting per unit length perpendicular to the tangential line drawn on the surface

of liquid. Ø Viscosity : It is a measure of resistance to flow which arises due to the internal friction between layers of fluid as

they slip past one another while liquid flows. Ø Coefficient of Viscosity : It is the force when velocity gradient (du/dz) and area of contact (A) both are unity.

Know the Facts Ø Liquids have definite volume due to presence of strong attractive intermolecular forces. Ø The liquids have indefinite shape. They take the shape of the container in which they are placed. Because the

molecules of liquids are in a state of constant random motion. This property of liquid is responsible for fluidity. Ø Liquids are much less compressible than gases because the intermolecular distances of separation are much

smaller in liquids than in gases. Ø At low atmospheric pressure, the liquid will boil at a lower temperature i.e. the boiling point decreases. So, it is

essential to use pressure cooker for cooking food at high altitudes. Ø At high external pressure, more heat is required to make the vapour pressure equal to external pressure i.e.

boiling point of the liquid increases. Ø The gases that decompose at their boiling temperature are purified by distillation under reduced pressure. Ø The S.I. unit of surface tension is Nm–1. Ø Sharp glass edges are heated to make them smooth. This is because on heating, the glass melts and takes up

rounded shape at the edges which has minimum surface area. This is known as fire polishing of glass. Ø Cleaning action of soap and detergents is due to lowering of interfacial tension between water and greasy

substances. Ø Surface tension decreases with increase in temperature. It is due to the fact that force acting per unit length

decreases due to increase in the kinetic energy of molecules. Ø Surface tension increases with increase in external pressure due to increase in force of attraction. Ø Surface tension is different for different liquids due to different intermolecular forces. Ø CGS unit of coefficient of viscosity = Poise Ø Poise in named after great scientist Jean Louise Poiseuille. Ø 1 poise = 1 g cm–1 = 10–1 Kg m–1 s–1.

qq

Chapter - 6 : Chemical Thermodynamics

TOPIC-1Types of System, Heat, Work, Energy and Enthalpy

Quick Review Ø Thermodynamics is the branch of science that deals with quantitative aspects of interconversion of various forms

of energy and the conversion of energy into work and vice versa. Ø System—Part of universe selected for observations. Ø Surrounding—Part of universe other than the observations.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 45 Ø System can be classified as— Open system can exchange matter and energy both with its surroundings. Closed system can exchange only energy with its surroundings but not the matter. Isolated system can neither exchange matter nor energy with surroundings. Ø State of system—The conditions of existence of a system when its macroscopic properties have a definite value. Ø Properties of a system can be classified into state and path function. Ø State Functions—The thermodynamic quantities which depend only on initial and final state of the system. e.g. pressure, volume and termperature are the variables which are used to describe the state of the system. Ø Path function—depends upon the path taken by the system during a change in state, e.g., heat and work. Ø Thermal process is the one which is carried out at a constant temperature. Ø Isobaric process is the one which is carried out at a constant pressure. Ø Isochoric process is the one during which the volume remains constant. Ø Adiabatic process is the one during which no heat is gained or lost by the system. Ø A reversible process takes place in a manner that system never deviates significantly from equilibrium state. It

occurs infinitesimally slowly. All real processes are irreversible in nature. Ø A cyclic process is the one that after undergoing a change in state returns back to its initial state. Ø An extensive property is a property whose value depends on the quantity or size of matter present in the system,

e.g., mass, volume, internal energy, enthalpy, heat capacity, etc. Ø An intensive property is that property which is independent of the mass or quantity or size of matter present in

the system e.g., temperature, density, pressure etc. Ø A substance (solid, liquid or gas) is said to be in standard state when it is pure, under 1 bar pressure and at any

specified temperature. Ø Heat and work are the two common forms of energy which are exchanged between a system and surroundings.

Both are path functions and appear only during a change in state. Ø Heat (q)—It is an exchange of energy which results in a temperature difference. Ø Work (w)—It is an organised form of energy.Ø Heat is positive if it is gained by the system (i.e., heat is transferred from the surroundings to the system) and heat

is negative if it is lost by the system (i.e., heat is transferred from system to the surroundings.) Ø Work is positive when it is done on the system and negative when done by the system. Ø Internal energy (U)—It is a type of energy associated with the system at particular condition of temperature and

pressure. Or Internal energy is the quantity that represents all the form of the enegy of the system, i.e., kinetic and potential

energies of the system. Ø Internal energy change (∆U)—It is a measure of heat change occur during the process at constant temperature

and constant volume.

∆U = qn

Ø Enthalpy (H)—It is the sum total of internal energy and PV energy of the system at particular conditions of temperature and pressure. It is also called heat content of the system.

H = U + PV.

Ø Enthalpy change (∆H)—It is the measure of heat change taking place during the process at constant temperature and constant pressure.

∆H = qp

∆H is negative for exothermic reactions and is positive for endothermic reactions. Ø Law of conservation of energy—It is also called first law of thermodynamics and states that energy of universe

always remains constant during physical or chemical changes.Or

Energy can neither be created nor be destroyed. Mathematically, it can be written as :

∆U = q + W

Ø Relation between ∆H and ∆U— ∆H = ∆U + P∆V

∆H = ∆U + ∆nRT

Ø The magnitude of temperature change during heat transfer depends upon the heat capacity of the system according to the relation.

q = m × c × ∆t Ø Specific heat or specific heat capacity is the quantity of heat required to raise the temperature of one unit mass of

a substance by one degree celsius (or one kelvin).


Know the Terms Ø Boundary : The wall that separates the system from the surroundings is called boundary. Ø Adiabatic System : The system which does not allow exchange of heat between the system and surroundings

through its boundary is called adiabatic system. Ø Reversible Process : A process is said to be reversible, if a change is brought out in such a way that the process

could, at any moment, be reversed by an infinitesimal change. Ø Free Expansion : Expansion of a gas in vacuum (Pex = 0) is called free expansion. Ø Exothermic Reactions : The reactions which evolve heat during the reaction are called exothermic reactions. Ø Endothermic Reactions : The reactions which absorb heat during the reaction are called endothermic reactions. Ø Extensive Property : It is a property whose value depends on the quantity or size of matter present in the system. Ø Intensive Property : It is a property whose value does not depend on the quantity or size of matter present n the

system. Ø Work : It is a mode of transfer of energy when system and surroundings differ in pressure till their pressures

become equal. Ø Temperature : It is a measure of degree of hotness or coldness. Ø Heat : The quantity of energy which flows between system and surroundings on account of temperature

difference is known as heat. Ø Heat Capacity : It is defined as the quantity of heat required to raise the temperature of a system by 1°C. Ø Specific Heat : It is defined as the quantity of heat required to raise the temperature of one unit mass by 1°C. Ø Heat Capacity at Constant Volume (Cv) : It is defined as the amount of heat required to raise the temperature of

one mole of a gas by one degree at constant volume of the gas. Ø Heat Capacity at Constant Pressure (Cp) : It is defined as the amount of heat required to raise the temperature of

one mole of a gas by one degree at constant pressure of the gas.

Know the Facts Ø Chemical energy stored by molecules can be released as heat during chemical reactions when a fuel burns in air. Ø The chemical energy may be used to do mechanical work when a fuel burns in an engine or to provide electrical

energy through a galvanic cell like dry cell. Ø The laws of thermodynamics deal with energy changes of macroscopic systems involving a large number of

molecules rather than microscopic systems containing a few molecules. Ø A system in thermodynamics refers to that part of universe in which observations are made and remaining

universe constitutes the surroundings. The Universe = The System + The Surroundings Ø Thermodynamics deals with matter in bulk and the properties arise due to bulk behavior of matter are called

macroscopic properties of the system. Ø The macroscopic properties of a gas i.e. pressure (P), volume (V), temperature (T), amount (n) etc. are called state

variables or state functions as their values depend only upon the initial and final states of system, but not on the path followed.

Ø Work done W = F × dx where F = force dx = distance Ø Work done is a path function not a state function as it depends upon the path followed. Ø Work done on the system increases the energy of the system and is considered as positive.

Work done on the system = Positive

Ø Work done by the system is considered as negative.

Work done by the system = Negative

Ø If a system loses energy, work is done by the system. Ø If a system gains energy, work is done on the system. Ø Thermodynamic Processes : l Isothermal process : ∆T = 0 l Adiabatic process : ∆q = 0 l Isobaric process : ∆P = 0 l Isochoric process : ∆V = 0 l Cyclic process : ∆U = 0

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 47 l Reversible Process : Process which proceeds infinitely slowly by a series of equilibrium steps. l Irreversible Process : Process which proceeds rapidly and the system does not have chance to achieve

equilibrium. Ø Sign Convections for Heat (Q) q = +ve, if heat is absorbed by the system q = –ve, if heat is evolved by the system Ø Work of expansion/compression : W = –Pext (Vf – Vi) where Pext = external pressure Vi = initial volume Vf = final volume Ø Work done in isothermal reversible expansion of an ideal gas :

Wrev = −2 303. lognRTVV

f

i

Or Wrev = − 2 303. lognRTPP

i

f

Ø Significance of ∆H and ∆U : ∆H = qp and ∆U = qv Ø Relation between ∆H and ∆U : For gaseous reaction, ∆H = ∆U + (np – nr)RT l If (np – nr) = 0, then ∆H = ∆U e.g. H2(g) + I2(g) → 2HI(g)

l If (np – nr) = positive, then ∆H > ∆U e.g. PCl5(g) → PCl3(g) + Cl2(g)

l If (np – nr) = negative, then ∆H < ∆U e.g. N2(g) + 3H2(g) → 2NH3(g) Ø q=C.∆T where C = Heat capacity ∆T=Changeintemperature q = amount of heat Or q = Cs×m×∆T where Cs = specific heat capacity m = mass of substance Or q = Cm × n×∆T where Cm = molar heat capacity n = No. of moles of a substance Ø According to first law of thermodynamics, ∆U=q + W where ∆U=(U2 – U1) = ∆E (change in internal energy) q = heat W = work done on the system Ø Whenasystemundergoesisothermalchange,∆U=0 i.e.thereisnoincreaseordecreaseintheinternalenergyofthesystem(∆U=0)thenfirstlawreducesto 0 = q + W or q = – W Ø When a system undergoes adiabatic change q = 0, i.e. there is no exchange of heat between system and

surrounding, the first law reduces to : ∆U=Wad ∴ The work is done at the expense of internal energy. Ø Limitations of First Law of Thermodynamics : l This law fails to explain the extent and direction of the convertibility of one form of energy into another. l This law fails to explain that how much energy would be transferred from one system to another. Ø We can measure energy changes associated with chemical or physical processes by an experimental technique

called calorimetry. Ø In calorimetry, the process is carried out in a vessel called calorimeter. Ø Cp – Cn = R


TOPIC-2Enthalpy Change

Quick Review Ø Enthalpy of Reaction (∆rH)—The enthalpy change accompanying the chemical reaction in which number of

moles of reactants consumed and those of products formed are the same as the stoichiometric coefficients. Ø Enthalpy of Neutralisation (∆nH)—It is the enthalpy change taking place during neutralisation of 1 gm equivalent

of acid with 1 gm equivalent of a base in dilute aqueous solution. Ø Standard Enthalpy of Fusion (∆fusionHΘ)—It is the enthalpy change taking place during the fusion of one mole

solid at its melting point. Ø Standard Enthalpy of Vaporisation (∆vapHΘ)—It is the enthalpy change taking place during the vaporisation of 1

mole of liquid at its boiling point and under standard pressure (1 bar). Ø Standard Enthalpy of Sublimation, (∆subHΘ)—It is the change in the enthalpy when one mole of a solid substance

sublimes at a constant temperature and under standard pressure (1 bar). Ø Standard Enthalpy of Formation (∆fH

Θ)—It is enthalpy change accompanying the formation of 1 mole of the substance from its constituent elements in their standard state.

∆fH° can be > 0 or < 0. Ø Factors affecting Enthalpy of Phase Transformation—The enthalpy change depends upon the strength of

intermolecular forces working on the substances that undergo phase transformation. Ø Hess’s law—The enthalpy change in a chemical or physical process is same whether the process is carried out in

one step or in several steps. Ø Applications of Hess’s Law: l to determine enthalpies of formation of compounds. l to determine enthalpies of extremely slow reactions. l to determine enthalpies of transformation of one allotropic form into another. l to determine bond energies (∆ Hreaction = S bond energies of reactants – S bond energies of products) l to determine resonance energy Ø Standard Enthalpy of Combustion (∆cH

Θ)—It is the enthalpy change occurring during the combustion of the mole of the substances in excess of oxygen. ∆cH° is always less than zero.

Ø Enthalpy of Atomisation (∆aHq)—It is the enthalpy change accompanying the dissociation of 1 mole of substance into gaseous atoms.

Ø Bond Enthalpy (∆bondHΘ)—The average amount of energy required to break one mole of the bonds of a particular type in gaseous molecules.

Ø Enthalpy of Solution (∆solHq)—It is the enthalpy change taking place when 1 mole of the solute is dissolved in large excess of solvent so that on further dilution no enthalpy occurs.

Ø Enthalpy of Hydration (∆hydHq)—It is the enthalpy change occurring during the hydration of 1 mole of anhydrous salt by combining with specific number of moles of water.

Ø Lattice Enthalpy (∆latticeHΘ)—Lattice Enthalpy of an ionic compound is the enthalpy change which occurs when one mole of an ionic compound dissociates into its ions in gaseous state.

Know the Terms Ø Thermochemical Equation : It is a balanced chemical equation in which the value of enthalpy of reaction (∆rH

Θ)is mentioned.

Ø Standard State : The most stable state of aggregation of an element at 25°C and 1 bar pressure is called its standard state. It is also called reference state.

Ø Calorific Value : It is the amount of heat produced on complete combustion of one gram of a fuel. Its unit is cal/g. Ø Bond Dissociation Enthalpy : It is defined as the enthalpy change accompanying the breaking of one mole of

covalent bond of a gaseous compound to give products in the gaseous phase. Ø Reaction Enthalpy : The enthalpy change accompanying reaction is called the reaction enthalpy.

Know the Facts Ø Negative ∆rH

Θ means the reaction is exothermic and positive ∆rHΘ means the reaction is endothermic.

Ø The unit of enthalpy change is kJ mol–1.

Ø On reverse a reaction, the sign of ∆H is also reversed but its value remains the same.

Ø Enthalpy of reaction ∆rH = Σ Σi i products i ia H b H− reactants

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 49 where S

i i productsa H = sum of enthalpies of products

Si i reac tsb H tan = sum of enthalpies of reactants

Ø The standard enthalpy of reaction (∆rHΘ) is the enthalpy change for a reaction when all the participating

substances are in their standard states. Ø H2O(s) → H2O(l) ; ∆fusH

Θ = 6.0 kJ mol–1

Ø H2O(l) → H2O(g) ; ∆vapHΘ = 40.79 kJ mol–1

Ø Sublimation is direct conversion of a solid into its vapour. Ø Solid CO2 or dry ice sublimes at 195 K with ∆subHΘ = 25.2 kJ mol–1. Ø Naphthalene sublimes slowly with ∆subHΘ = 73.0 kJ mol–1. Ø The magnitude of the enthalpy change depends on the strength of the intermolecular interactions in the substance

undergoing the phase transformations. Ø For acetone, intermolecular dipole-dipole interactions are significantly weaker. So, it requires less heat to vaporize

1 mol of acetone than it does to vaporize 1 mol of water. Ø Some reactions with standard molar enthalpies of formation are given below : • H2(g) + ½O2(g) H2O(l); ∆fH

Θ = – 285.8 kJ mol–1

• C(graphite,s)+2H2(g) → CH4(g) ; ∆fHΘ = –74.81 kJ mol–1

• 2C(graphite,s)+3H2(g) + ½ O2(g) → C2H5OH(l) ; ∆fHΘ = 277.7 kJ mol–1

Ø Decomposition of CaCO3(s) is an endothermic process. CaCO3(s) → CaO(s) + CO2(g); ∆rH

Θ= 178.3 kJ mol–1

Ø C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l); ∆rHΘ = – 1367 kJ mol–1

This reaction is exothermic as it possesses negative value of ∆rHΘ

Ø Combustion reactions are exothermic in nature. Ø Cooking gas in cylinders contains mostly butane (C4H10). During complete combustion of one mole of butane,

2658 kJ of heat is released.

C4H10(g) + 132

O2(g) → 4CO2(g) + 5H2O(l); ∆cHΘ = – 2658 kJ mol–1

Ø Combustion of glucose gives out 2802.0 kJ mol–1 of heat. C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l); ∆cH

Θ =– 2802 kJ mol–1

Ø Our body generates energy from food by the overall process same as combustion of glucose, although the final products are produced after a series of complex bio-chemical reactions involving enzymes.

Ø Examples of enthalpy of atomization : CH4(g) → C(g) + 4H(g); ∆aHΘ = 1665 kJ mol–1

Na(s) → Na(g); ∆aHΘ = 108.4 kJ mol–1

TOPIC-3Spontaneity

Quick Review Ø Spontaneous process—A process which has an urge or a natural tendency to occur in a particular direction either

of its own or after proper initiation under a given set of conditions. Ø Entropy is a thermodynamic property which is a measure of disorder. It is related to heat absorbed at a constant

temperature by the system in a reversible process as :

∆S = qrev

T For a spontaneous change, total entropy change is positive. Ø Second law of Thermodynamics : The entropy of universe always tends to increase during any spontaneous

process. Ø Third law of Thermodynamics : At absolute zero, the entropy of a perfectly crystalline substance is zero. Ø Gibbs energy is another thermodynamic property which is a measure of energy stored in a system or a substance

that is available for doing useful work. It is related to changes in enthalpy and entropy of the system during a process as :

∆G = ∆H – T∆S For a spontaneous change, ∆Gsys < 0, for a non-spontaneous change, ∆Gsys > 0, and ∆Gsys = 0 shows equilibrium.


Important Formulae (i) ∆U = q + W (ii) ∆H = ∆U + P∆V (iii) ∆H0 = S∆fH

0 (Products) – S∆fH° (Reactants) (iv) ∆H0 = [Sum of bond energies of reactants] – [Sum of bond energies of products]

(v) ∆S = q rev( )

T (vi) For process to be spontaneous ∆Ssystem + ∆Ssurr > 0 (vii) ∆G = ∆H – T∆S (viii) ∆Gsys = – T∆STotal

(ix) ∆rG° = S∆fG° (Products) – SfG° (Reactants)

(x) ∆GT, P < 0 refers to spontaneous process.

Know the Terms Ø Entropy : It is the measure of degree of randomness or disorderness in an isolated system. Ø Phase Transition : The change of matter from one state to another is known as phase transition. Ø Gibbs Energy : It is the maximum amount of energy available to a system, during a process, that can be converted

into useful work. It is also called free energy or Gibbs free energy. Ø Standard Free Energy Change : It is defined as the free energy change for a process at 298K and 1 atm pressure

in which the reactants in their standard state are converted into products in their standard state. It is denoted by ∆G°.

Ø Standard Free Energy of Formation : It is defined as the free energy change when 1 mole of the compound is made from its elements in their standard states.

Know the Facts Ø The first law of thermodynamics tells us about the relationship between the heat absorbed and the work

performed on or by a system. It puts no restrictions on the direction of heat flow. Ø The flow of heat is unidirectional from higher temperature to lower temperature. Ø All naturally occurring processes whether chemical or physical will tend to proceed spontaneously in one

direction only. For example, a gas expanding to fill the available volume, burning carbon in dioxygen giving carbon dioxide etc.

Ø A spontaneous process is an irreversible process and may only be reversed by some external energy. Ø A system at higher temperature has greater randomness in it than one at lower temperature.

Ø Entropychange∆S=q

Trev

Ø ∆Stotal = ∆Ssystem + ∆Ssurr > 0 Ø When a system is in equilibrium, the entropy is maximum and the change in entropy, ∆S = 0 Ø The entropy for a spontaneous process increases till it reaches maximum and at equilibrium the change in entropy

is zero. Since entropy is a state property, we can calculate the change in entropy of a reversible process by :

∆Ssys = q

Tsys rev.

Ø For reversible and irreversible expansion for an ideal gas, under isothermal conditions, ∆U = 0, but ∆Stotal i.e. ∆Ssys +∆Ssurr is not zero for irreversible process. Hence ∆U does not discriminate between reversible and irreversible process, whereas ∆S does.

Ø G = H – T S Gibbs function, G is an extensive property and a state function. Ø ∆G = ∆H – T∆S (Gibbs Helmholtz equation) Ø If the system is in thermal equilibrium with the surrounding, then the temperature of the surrounding is same as

that of the system. Increase in enthalpy of the surrounding is equal to decrease in the enthalpy of the system. Ø For spontaneous process ∆Stotal > U Ø ∆G (Change in free energy) gives a criteria for spontaneity at constant pressure and temperature. (i) If ∆G is negative (< 0), the process is spontaneous. (ii) If ∆G is positive (> 0), the process is non-spontaneous.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 51 Ø If a reaction has a positive enthalpy change and positive entropy change, it can be spontaneous when T∆S is large

enough to out weigh ∆H. This can happen in two ways : (i) The positive entropy change of the system can be small in which case ‘T’ most be large. (ii) The positive entropy change of the system can be ‘large’ in which case ‘T’ may be small. Ø ∆rG

Θ = – RT lnK = – 2.303 RT log K where R = Gas constant T = Temperature (in Kelvin) K = Equilibrium constant Ø If ∆G = 0, the process is in equilibrium.

qq

Chapter - 7 : Equilibrium

TOPIC-1Physical and Chemical Equilibrium

Quick Review Ø Reversible Reaction : A chemical reaction which can take place in forward as well as backward direction. Ø Irreversible Reaction : A chemical reaction which cannot take place in the backward direction. Ø Equilibrium State : A state of the system when its observable properties do not change at a given set of conditions.

At equilibrium, the forward and backward reactions proceed at same rate. Ø The state of equilibrium in which the rate of a forward reaction and rate of backward reaction are equal and there

is no net change in composition is known as dynamic equilibrium. Ø Equilibrium reached in physical processess is termed as physical equilibrium. It may involve an equilibrium

between different phases like solid-liquid, liquid-vapour and solid-vapour phases or a dissolution equilibrium of a solid or gaseous solute in a liquid solvent.

Ø Same state of equilibrium can be achieved by taking the reactant or product or a mixture of both. Ø Chemical Equilibrium : Equilibrium achieved in chemical processes. Ø Law of Mass Action : A theoretical concept which states that rate of reaction is directly proportional to the product

of molar concentrations of reactants raised to the power equal to stoichiometric coefficient of that species in the balanced equation.

Ø Rate constant : It is the rate of reaction, when concentration of each reacting species is unity. Ø The state of chemical equilibrium is characterised by its equilibrium constant Kc or Kp. Ø Kc , concentration equilibrium constant is calculated by using equilibrium concentrations of all the reactants and

products. Ø Kp , pressure equilibrium constant is calculated by using the equilibrium partial pressures of gaseous reactants

and products. Ø Equilibrium constant : It is the ratio of the rate constants of forward reaction to that of backward reaction. Ø Law of Chemical Equilibrium : At a given temperature, the product of concentrations of the reaction products

raised to the respective stoichiometric coefficient in the balanced chemical equation divided by the product of concentrations of the reactants raised to their individual stoichiometric coefficients has a constant value. This is known as Law of Chemical Equilibrium

For an equilibrium,

aA bB cC+ +� dD

Equilibrium constant (Kc) is expressed as –

Kc = C D

A B

c d

a b[ ] [ ][ ] [ ]

where [A] and [B] = Equilibrium concentrations of the reactants [C] and [D] = Equilibrium concentrations of the products Ø Equilibrium Constant in Gaseous Systems: Consider a general reaction,

aA bB cC dD+ +�


Kp = P P

p pC

cD

d

Aa

Bb

[ ] [ ][ ] [ ]

= C D RT

A B RT

c d c d

a b a b

[ ] [ ] ( )[ ] [ ] ( )

+( )

+( )

= C D

A BRT

c d

a bc d a b[ ] [ ]

[ ] [ ]+ − +( )( ) ( )

= C D

A BRT

c d

a bn[ ] [ ]

[ ] [ ]∆( )

Kp = Kc (RT)∆n

where R = Gas constant T = Absolute temperature ∆n = (Number of moles of gaseous products – Number of moles of gaseous reactants) in balanced

chemical equation. = (c + d) – (a + b) Ø Homogeneous Equilibria : In a homogeneous system, all the reactants and products are in the same phase. For example, in the gaseous reaction, below reactants and products are in homogeneous phase.

N g H2 23( ) (g) 2NH (g)3+ �

Ø Heterogeneous Equilibria : Equilibrium in a system having more than one phase is called heterogeneous equilibrium.

For example, equilibrium between water vapour and liquid water in a closed container

H O2 ( ) H O(g)2l �

Ø Applications of Equilibrium Constant : To predict the extent of a reaction on the basis of its magnitude. •To predict the direction of the reaction. •To calculate equilibrium concentrations. Ø From the value of equilibrium constant, we can predict the extent of reaction. If Kc > 103 that means products

predominate over reactants and reaction proceeds nearly to completion. If Kc < 10–3, reactants predominate over products, i.e., reaction proceeds rarely. If Kc is in the range of 10–3 to 103, concentrations of both reactants and products are in appreciable amount.

Ø Reaction Quotient (Q) : The ratio of the product of molar concentrations of the products to the product of molar concentrations of reactants with each concentration terms raised to the power equal to the stoichiometric coefficient of that species is called reaction quotient.

At equilibrium Q = K. At concentration other than equilibrium Q ≠ K. If Qc > Kc, the reaction, will proceed in the direction of reactants (backward reaction). If Qc < Kc, the reaction will proceed in the direction of the products (forward reaction). If Qc = Kc, the reaction mixture is already at equilibrium. Ø K = e –∆G°/RT

From this equation, we can predict the spontaneity of the reaction. If ∆G° < 0, then – ∆G°/RT is positive and e –∆G°/RT > 1, ∴ K > 1, the reaction is spontaneous. If ∆G° > 0, then – ∆G°/RT is negative and e –∆G°/RT < 1, ∴ K < 1, the reaction is non-spontaneous. Ø ∆G = ∆G° + RT ln Q At equilibrium, ∆G° = – RT ln K = – 2·303 RT log K. Ø Le Chatelier’s Principle : When an equilibrium is subjected to any kind of stress (change in concentration,

temperature or pressure) it shifts in a direction so as to undo the effect of stress. Ø Increase in concentration of any reactant shifts the equilibrium in the forward direction, i.e., net reaction occurs in

the forward direction. However, on decreasing the concentration of any reactant, the reaction moves in backward direction. Similar changes in concentration of any product result in opposite changes.

Ø On increasing the pressure, the reaction occurs in that direction in which number of moles of the gas or pressure decreases or the direction in which ∆ng is negative and vice-versa.

Ø If the reaction involves solids and/or liquids, i.e., ∆ng = 0, then there is no effect of pressure change on the system. Ø The equilibrium constant for an exothermic reaction decreases as the temperature increases.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 53 Ø The equilibrium constant for an endothermic reaction increases as the temperature increases. Ø Addition of inert gas at constant volume does not affect the equilibrium. Ø Addition of catalyst does not alter the equilibrium, it just help in attaining the equilibrium faster.

Know the Terms Ø Equilibrium—It is a state in a process when two opposing processes (forward and backward) occur simultaneously

at the same rate. Ø Equilibrium Mixture—The mixture of reactants and products in the equilibrium state is called an equilibrium

mixture. Ø Dynamic Equilibrium—The stage at which the rates of the forward and backward reactions become equal is

called dynamic equilibrium. Ø Ionic Equilibrium—The equilibrium which involve the ions in aqueous solutions is called ionic equilibrium. Ø Normal Melting Point or Normal Freezing Point—For any pure substance at atmospheric pressure, the

temperature at which the solid and liquid phases are at equilibrium is called the normal melting point or normal freezing point of the substance.

Ø Henry’s Law—It states that the mass of a gas dissolved in a given mass of a solvent at any temperature is proportional to the pressure of the gas above the solvent.

m a P

Know the Facts Ø Chemical equilibria are important in numerous biological and environmental processes. For example, equilibria involving O2 molecules and the protein haemoglobin play an important role in the

transport and delivery of O2 from our lungs to our muscles. Ø Equilibrium can be established for both physical processes and chemical reactions. Ø The reaction may be fast or slow depending on the experimental conditions and the nature of reactants. Ø The boiling point of water is 100°C at 1.013 bar pressure. Ø For any pure liquid at one atmospheric pressure (1.013 bar), the temperature at which the liquid and vapours are

at equilibrium is called normal boiling point of the liquid. Ø Boiling point of the liquid depends on the atmospheric pressure. At high altitude, the boiling point decreases. Ø For solid � liquid equilibrium, there is only one temperature (melting point) at 1 atm (1.013 bar) at which the

two phases can coexist. If there is no exchange of heat with the surroundings, the mass of the two phases remains constant.

Ø For liquid � vapour equilibrium, the vapour pressure is constant at a given temperature. Ø For dissolution of solids in liquids, the solubility is constant at a given temperature. Ø For dissolution of gases in liquids, the concentration of a gas in liquid is proportional to the pressure (concentration)

of the gas over the liquid. Ø Some Features of Physical Equilibria – l Process Conclusion Liquid � Vapour PH2O constant at given temperature. H2O(l) � H2O(g) l Solid � Liquid Melting point is fixed at constant pressure H2O(S) � H2O(l) l Solute(s) � Solute (solution) Concentration of solute in solution is constant at a given temperature Sugar(s) � Sugar (solution)

l Gas(g) � Gas(aq) [gas(aq)]/[gas(g)] is constant at a given temperature CO2(g) � CO2(aq)

Ø Relation between Equilibrium Constants for a General Reaction and its Multiples—

Chemical Equation Equilibrium Constant

l aA + bB � cC + dD K

l cC + dD � aA +bB Kc’ = (1/Kc)

l naA + nbB � ncC + ndD Kc” = (Kcn )

l 1/n aA + 1/n bB � 1/ncC + 1/ndD Kc” = (Kc1/n)

54 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø 1 pascal (Pa) = 1 Nm–2

1 bar = 105 Pa Ø The free energy change at equilibrium state is zero i.e. ∆G = 0 Ø The value of the equilibrium constant is not affected by the addition of a catalyst to the reaction, it is due to the

fact that the catalyst increases the rate of the forward reaction and the backward reaction to the same extent. Ø The value of the equilibrium constant changes with change to temperature.

log

.KK

HR T T

2

1 1 22 3031 1

=∆

−

Ø The value of equilibrium constant is independent of initial concentration of the reactants.

TOPIC-2lonic Equilibrium: Acids, Bases and their lonization

Quick Review ØSubstances that conduct electricity in aqueous solutions are called electrolytes, e.g., acids, bases and salts are

electrolytes and their solution conduct electricity due to the process of ions formed by dissociation of electrolytes in aqueous solution.

ØStrong electrolytes completely dissociate in their solutions.ØWeak electrolytes dissociate partially and an equilibrium exist between the ions and the undissociated electrolyte.ØAcid—A substance which furnishes H+ ions in aqueous solution (Arrehenius concept); proton donor (Bronsted

concept) and acceptor of electron pair (Lewis concept).ØBase—A substance which gives OH– ions in aqueous solution (Arrehenius concept); acceptor of proton (Bronsted

concept) and donor of electron pair (Lewis concept).ØNeutralisation—A combination of H+ and OH– ions in solution (Arrhenius concept); transferance to proton

(Bronsted concept); formation of dative bond (Lewis concept).ØThe acid-base pair that differs only by one proton is called a conjugate acid-base pair. If Bronsted acid is a strong

acid then its conjugate base is a weak base and vice-versa.ØIonic product of water, Kw , is the product of the concentration of hydronium ions and hydroxyl ions in pure

water at a particular temperature. Kw = [H3O+][OH]–

= 1 × 10–14 mol2 L–2 at 298 KØpH+ of solution—Potency of H+ in solution. It is negative logarithm of H3O+ ion concentration in solution.ØpOH– of solution—It is negative logarithm of OH– ion concentration of the solution.ØFor any solution at 25°C. pKw = pH + pOH = 14 pH = – log [H3O+] pOH = – log [OH–]; pKw = – log Kw

ØKw ionic product of water (H+)(OH–) = 1 × 10–14 at 25°C.ØStrength of acids or bases is measured in terms of their respective dissociation constants Ka (or pKa) and Kb or

(pKb).ØLarger the value of Ka or lower value of pKa corresponds to greater acidic strength of acids. Similarly larger the

value of Kb or lower value of pKb corresponds to more stronger base.ØDissociation constants of conjugate acid-base pair to each other as : Ka.Kb = Kw

pKa + pKb = pKw = 14 (at 298 K)ØFor monobasic acid HA, if C is the molar concentration of HA and a is the degree of dissociation then

(H3O+) = Ca = K Ca × Ø Similarly of monoacid base [OH–] = Ca = K Cb ×ØMany acids and bases ionise in two or more steps; These have separate ionisation constant for each step of

ionisation. Strength of an acid depends upon (i) strength of H—A bond : weaker the bond, stronger is the acid and (ii)

polarity of H—A bond : greater the polarity, stronger is the acid.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 55ØDissociation of weak acids and bases is suppressed in the presence of common ion furnished by their salts. This

phenomenon is known as common ion.ØHydrolysis : It is reverse of neutralization. It involves the interactions of ions of electrolyte with H2O molecules

in solution to give acidic or basic solution.ØSalts of strong acids and strong bases do not undergo hydrolysis and their aqueous solutions are neutral.ØSalts of strong acids with weak bases undergo cationic hydrolysis and give acidic solutions.ØSalts of strong bases with weak acids undergo anionic hydrolysis and give basic aqueous solutions.

Know the Terms Ø Electrolytes : The substances that conduct electricity either in molten state or in their aqueous solutions are called

electrolytes. Ø Strong Electrolytes : The substances which dissociate completely in to ions in aqueous solution are called strong

electrolytes. Ø Weak Electrolytes : The substances which do not dissociate completely into ions in aqueous solution are called

weak electrolytes. Ø Ionic Equilibrium : It is defined as the equilibrium established between the unionized molecules and the ions in

a solution of weak electrolytes. Ø Arrhenius Acid : It is the substance that dissociates in water to give hydrogen ions (H+). Ø Arrhenius Base : It is the substance that dissociates in water to give hydroxyl ions (OH–). Ø Bronsted – Lowry Acid : It is a substance that is capable of donating a proton (H+) Ø Bronsted Lowry Base : It is a substance that is capable of accepting a proton (H+) Ø Conjugate Acid – Base Pair : The acid base-pair that differs only by one proton is called a conjugate acid-base pair. Ø Lewis Acid : It is defined as a species which accepts electron pair. Ø Lewis Base : It is defined as a species which donates an electron pair. Ø pH Scale : Hydronium ion concentration in molarity is more conveniently expressed on a logarithmic scale

known as the pH scale. Ø pH : The pH of a solution is defined as the negative logarithm to base 10 of the activity (aH

+) of hydrogen ion. Ø Acid Ionization Constant : The equilibrium constant for acid ionization is called acid ionization constant. Ø Base Ionization Constant : The equilibrium constant for base ionization is called base ionization constant. Ø Common Ion Effect : It can be defined as a shift in equilibrium on adding a substance that provides more of an

ionic species already present in the dissociation equilibrium. Ø Hydrolysis : The process of interaction between water and cations or anions or both of salts is called hydrolysis.

Know the Facts Ø Hydrochloric acid present in the gastric juice is secreted by the lining of our stomach in a significant amount of

1.2 – 1.5 L/day and is essential for digestive process. Ø Acetic acid is known to be the main constituent of vinegar. Ø Lemon and orange juices contain citric acid and ascorbic acids. Ø Tartaric acid is found in tamarind paste. Ø As most of the acids taste sour, the term `acid’ has been derived from a Latin word “acidus” meaning sour. Ø Acids turn blue litmus paper into red. Ø Acids liberate dihydrogen on reacting with some metals. Ø Bases turn red litmus paper into blue. Ø Bases taste bitter and feel soapy. Ø A common example of a base is washing soda used for washing purposes. Ø When acids and bases are mixed in the right proportion, they react with each other to form salts. Ø The electrostatic forces between two charges are inversely proportional to dielectric constant of the medium. Ø Water acts as a universal solvent. It possesses a very high dielectric constant of 80. Ø When sodium chloride is dissolved in water, the electrostatic interactions are reduced by a factor of 80 and this

facilitates the ions to move freely in the solution. Ø The extent to which ionization occurs, depends upon the strength of the bond and the extent of solvation of ions

produced. Ø Dissociation refers to the process of separation of ions in water already existing as such in the solid state of the

solute, as in sodium chloride. While ionization corresponds to a process in which a neutral molecule splits into charged ions in the solution.

Ø A bare proton (H+) is very reactive and can not exist freely in aqueous solutions. So, it bonds to the oxygen atom of a solvent water molecule to give hydronium ion (H3O+). It’s geometry is trigonal bipyramidal.

56 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø If Bronsted acid is a strong acid then its conjugate base in a weak base and vice-versa. Ø The density of pure water is 1000 g/L and its molar mass is 18.0 g/mol.

Molarity of pure water = Density

Molar Mass

= 1000

18 0g/L

g/mol. = 55.55 M Ø We can distinguish acidic, neutral and basic aqueous solutions by the relative values of the H3O+ and OH–

concentrations : Acidic : [H3O+] > [OH–] Neutral : [H3O+] = [OH–] Basic : [H3O+] < [OH–] Ø pH range : 1 7 14 (Acidic solution) (pH<7) (Neutral solution) (pH = 7) (Basic solution) (pH>7) Ø pH = –log10 [H+] Ø [H+] = 10–pH

Ø Kw = [H3O+][OH–] = 10–14

where Kw = Ionic product of water Ø pKw = pH + pOH = 14 Ø Relative order of acidic strength among halogen acids : Size increases HF << HCl << HBr << HI Acid strength increases As the size of halogen (X) increases down the group, H – X bond strength decreases and therefore the acid

strength increases. Ø Relative order of acidic strength among CH4, NH3, H2O and HF (HA) is: electronegativity of ‘A’ increases CH4 < NH3 <H2O < HF Acid strength increases Ø Salt of weak acid and strong base: CH3COONa Ø Salt of strong acid and weak base: NH4Cl Ø Salt of weak acid and weak base: CH3COONH4

TOPIC-3Buffer Solution and Solubility Product

Quick Review Ø Buffer solution is generally a mixutre of weak acid and its conjugate base or weak base and its conjugate acid. Ø pH of buffer solution is given by Hendersons-Hasselbalch equation :

pH = pKa + log [ ][ ]

SaltAcid

(for acidic buffers)

pOH = pKb + log [ ][Base]

Salt (for basic buffers)

Ø pH = pKa when [Salt] = [Acid]

Ø pOH = pKb when [Salt] = [Base] Ø The dissolving of the sparingly soluble salt in water to form a saturated solution, equilibrium is established

between undissolved solid and the ions in a saturated solution. AB (s) �An+ (aq) + Bn– (aq) Ksp = [An+] [Bn–] = (S) (S) = S2

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 57 Ø Solubility Product (Ksp)—It is a product of the concentration of ions of electrolyte at the saturation point with each

concentration term raised to the power equal to numerical coefficient of that species in the balanced equation. It is the highest value of ionic product for a sparingly soluble salt.

Ø The solubility of a sparingly soluble salt further decreases due to Common Ion Effect. Ø Solubility of a salt of a weak acid increases as the pH of the solution decreases as the concentration of H3O+ ions

increases. Ø From the value of Ksp (solubility product) and Q (ionic product), we can predict whether the precipitation of salt

or dissolution of salt will take place. If Q > Ksp , then precipitation will take place. If Q < Ksp , then dissolution will take place. If Q = Ksp , then reaction is at equilibrium.

Know the Terms Ø Buffer Solution : The solutions which resist change in pH on dilution or with the addition of small amounts of

acid or with the addition of small amounts of acid or alkali are called Buffer Solution. Ø Solubility Product : It is a mathematical product of its dissolved ion concentrations raised to the power of their

stoichiometric coefficients.

Know the Facts Ø Many body fluids e.g. blood or urine have definite pH and any deviation in their pH indicates malfunctioning of

the body. Ø Buffer solutions of known pH can be prepared from the knowledge of pKa of the acid or pKb of the base and by

controlling the ratio of the salt and acid or salt and base. Ø A mixture of acetic acid and sodium acetate acts as buffer solution around pH 4.75. Ø A mixture of ammonium chloride and ammonium hydroxide acts as a buffer solution around pH 9.25. Ø For a salt to dissolve in a solvent, the strong forces of attraction between its ions (lattice enthalpy) must be

overcome by the ion-solvent interactions. Ø The solvation enthalpy of ions is referred to in terms of solvation which is always negative i.e. energy is released

in the process of solvation. Ø The amount of solvation enthalpy depends on the nature of solvent. If solvent is non-polar, the solvation enthalpy

is small and therefore, not sufficient to overcome lattice enthalpy of the salt.`qq

Chapter - 8 : Redox Reactions

TOPIC-1Concept of Oxidation-Reduction, Redox Reactions and Oxidation Number

Quick Review Ø Oxidation : According to classic concept, “it is defined as the addition of oxygen/electronegative element to a

substance or removal of hydrogen/electropositive element from a substance”. eg. CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) (Addition of oxygen) Mg(s) + F2(g) → MgF2(s) (Addition of electronegative element) 2H2S(g) + O2(g) → 2S(s) + 2H2O(l) (Removal of hydrogen) 2K4[Fe(CN)6](aq) + H2O2(aq) → 2K3[Fe(CN)6](aq) + 2KOH(aq) (Removal of electropositive element) Ø Reduction : According to classical concept, it is defined as the removal of oxygen/electronegative element from a

substance or addition of hydrogen/electropositive element to a substance. e.g.

2HgO(s) ∆ → 2Hg(l) + O2(g) (Removal of oxygen) 2FeCl3(aq) + H2(g) → 2FeCl2(aq) + 2HCl(aq) (Removal of electronegative element) CH2 = CH2(g) + H2(g) → CH3 – CH3(g) (Addition of hydrogen) 2HgCl2(aq) + SnCl2(aq) → Hg2Cl2(s) + SnCl4(aq) (Addition of electropositive element)

58 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Redox Reaction : The reaction which involves oxidation and reduction reactions simultaneously is called redox

reaction. e.g. 2 Nas(s) + Cl2(g) → 2NaCl(s) Ø Redox Reaction in Terms of Electron Transfer Reaction : Consider a redox reaction, 2Na(s) + Cl2(g) → 2NaCl(s) Na(s) → Na+ + e– (Oxidation half reaction) Cl2(g) + 2e– → 2Cl–(g) (Reduction half reaction) According to electronic concept, The half reaction that involves loss of electron is called oxidation half reaction and the half reaction that involves

gain of electron is called reduction half reaction. Ø Oxidising Agent : The species which can oxidise other substance and itself reduces is called oxidising agent. It is

acceptor of electrons. Ø Reducing Agent : The species which can reduce other substance and itself oxidises is called reducing agent. It is

donor of electrons. Ø Oxidation Number : Oxidation number denotes the oxidation state of an element in a compound ascertained

according to a set of rules formulated on the basis that electron pair in a covalent bond belongs entirely to more electronegative element.

Ø According to oxidation number concept, l Oxidation is defined as an increase in the oxidation number of the element in the given substance. l Reduction is defined as a decrease in the oxidation number of the element in the given substance. l Oxidising agent is a reagent which can increase the oxidation number of an element in a given substance. l Reducing agent is a reagent which can decrease the oxidation number of an element in a given substance. l Redox reactions are those reactions which involve change in oxidation number of the interacting species.

Know the Terms Ø Oxidation : l According to classical concept, it is defined as the process that involves addition of oxygen or other

electronegative element or removal of hydrogen or removal of other electropositive element. l According to electronic concept, it is defined as the process that involves loss of electrons. l According to oxidation number concept, it is defined as the process which involves increase in oxidation

number of the element in the given substance. Ø Reduction : l According to classical concept, it is defined as the process that involves addition of hydrogen or addition of

other electropositive element or removal of oxygen or removal of other electronegative element. l According to electronic concept, it is defined as the process that involves gain of electrons. l According to oxidation number concept, it is defined as the process which involves decrease in oxidation

number of the element in the given substance. Ø Oxidation number : It is defined as the charge that an atom of an element has on its ion or appear to have when

present in the combined state with other atoms.

Know the Facts Ø Transformation of matter from one kind into another occurs through the various types of reactions such as redox

reactions. Ø Redox reactions find extensive use in pharmaceutical, biological, industrial, metallurgical and agricultural areas. Ø The reaction between metallic zinc and the aqueous solution of copper nitrate is :

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) (Redox Reaction)

Zn(s) → Zn2+(aq) + 2e– (Oxidation half reaction)

Cu2+(aq) + 2e– → Cu(s) (Reduction half reaction) Ø Electron releasing tendency among Zn, Ag and Cu is in the order : Zn > Cu > Ag. Ø An element, in the free or the uncombined state, bears an oxidation number of zero. e.g. Each atom in H2, O2, O3, P4, Na, Mg etc. has the oxidation number zero. Ø For ions composed of only one atom, the oxidation number is equal to the charge on the ion. e.g. Ion Oxidation Number Na+ +1 Mg2+ +2 Fe3+ +3 Cl– – 1 O2– – 2

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 59 Ø The oxidation number of oxygen in most compounds is –2. Exception : (i) In case of peroxides and superoxides, the oxidation numbers of oxygen are –1 and –1/2 respectively. e.g. Peroxides = H2O2, Na2O2 (Oxidation number of O = –1) Superoxides = KO2, RbO2 (Oxidation number of O = –1/2) (ii) When oxygen is bonded to fluorine, then in such compounds e.g. oxygen difluoride (OF2) and dioxygen

difluoride (O2F2), the oxidation numbers of oxygen are +2 and +1 respectively. Ø The oxidation number of hydrogen in most compounds is +1. Exception : When hydrogen is bonded to metals in binary compounds. e.g. LiH, NaH and CaH2 , in such cases,

the oxidation number of hydrogen is –1. Ø The oxidation number of fluorine in all of its compounds is –1. Other halogens like Cl, Br and I also possess an oxidation number of –1, when they occur as halide ions in their

compounds. However, when Cl, Br or I is combined with oxygen like oxoacids and oxoanions, then these halogens possess

positive oxidation numbers. e.g. Cl2O7, in this compound, oxidation number of Cl is +7. Ø The algebraic sum of the oxidation number of all the atoms in a compound must be zero. Ø In poly atomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must equal the charge on

the ion. Ø A term that is often used inter changeably with the oxidation number is the oxidation state. e.g. In CO2, the oxidation state of carbon is +4, that is also its oxidation number. Ø The oxidation number denotes the oxidation state of an element in a compound. Ø The oxidation number or state of a metal in a compound is sometimes presented according to the notation given

by German chemist, Alfred Stock. It is popularly known as Stock notation. According to this notation, the oxidation number is expressed by putting a Roman numeral representing the

oxidation number in parenthesis after the symbol of the metal in the molecular formula. e.g. Stannous chloride in written as Sn(II)Cl2, Stannic chloride is written as Sn(IV)Cl4 etc.

TOPIC-2Types of Redox Reactions, Balancing of Redox Reactions and Electrode Processes

Quick Review Ø Redox reactions are of various types : 1. Combination Reactions : The reactions which involve combination of two or more atoms or molecules to

give only a single compound are called combination reactions.

e.g. C( ) ( ) ( )0 0

2 2

4 2

s O g CO g+ →+ −

∆

2. Decomposition Reactions : The reactions which lead to the breakdown of a compound into two or more components at least one of which must be in the elemental state are called decomposition reactions.

e.g. 2 21 1 0 0

2Na H s Na s H g+ −

→ +( ) ( ) ( )∆

3. Displacement Reactions : The reactions which involve replacement of an atom (or ion) by an atom (or ion) of another element from a compound are called displacement reactions. These reactions can be represented as:

e.g. X + YZ → XZ + Y These reactions can further be classified into following categories : (i) Metal Displacement Reactions : In these reactions, a metal in a compound can be displaced by another

metal in the uncombined state.

e.g. CuSO aq Zn s Cu s ZnSO aq4

2 6 2 0 0

4

2 6 2+ + + − + + −

+ → +( ) ( ) ( ) ( ) (ii) Non-metal Displacement Reactions : These reactions include hydrogen displacement and a rarely

occurring reaction involving oxygen displacement.

e.g. 2 2 20 1

2

2 1 2 1 0

2Na s H O l NaOH aq H g( ) ( ) ( ) ( )+ → ++ − + − +

Mg s H O g Mg OH g H g0 1

2

2 2 2 1

2

0

22( ) ( ) ( ) ( ) ( )+ → ++ − + − +


4. Disproportionation Reactions : These reactions are special type of redox reactions. In these reactions, an element in one oxidation state is simultaneously oxidized and reduced.

e.g. 2 21

2

1

2

1

2

2 0

2H O aq H O l O g+ − + −

→ +( ) ( ) ( )

P s OH aq H O l P H g H PO0

4 2

3

3 2 2

1

3 3 3( ) ( ) ( ) ( )+ + → +−−

−+

Ø Balancing of Redox Reactions : Following methods are used to balance chemical equations for redox processes : 1. Oxidation Number Method : It involves following steps : l Write the correct chemical formula for each reactant and product. l Identify atoms which undergo change in oxidation number in the reaction by assigning oxidation

number to all elements in the reaction. l Calculate the increase or decrease in the oxidation number per atom and for the entire molecule/ion in

which it occurs. If these are not equal, then multiply by suitable number so that these become equal. l Ascertain the involvement of ions if the reaction is occurring in water, add H+ or OH– to the expression

on the appropriate side so that the total ionic charges of reactants and products are equal. If the reaction is carried out in acidic solution, use H+ ions in the equation; if in basic solution, use OH–

ions. l Make the numbers of H–atoms in the expression on the two sides equal by adding H2O molecules to

the reactants or products. Now, check the number of O–atoms. If there are the same number of oxygen atoms in the reactants and

products, then the equation represents the balanced redox reaction. 2. Half Reaction Method : It involves following steps : l Produce unbalanced equation for the reaction in ionic form. l Separate the equation into half reactions i.e. oxidation and reduction half reactions. l Balance the atoms other than O and H in each half reaction individually. l For reactions occurring in acidic medium, add H2O to balance O atoms and H+ to balance H atoms. l Add electrons to one side of the half reaction to balance the charges. If required, make the number

of electrons equal in the two half reactions by multiplying one or both half reactions by appropriate number.

l Now, add two half reactions to achieve overall reaction and cancel the electrons on each side. This gives the net ionic equation.

l Verify that the equation contains the same type and number of atoms and the same charges on both sides of the equation.

Note: For the reaction in a basic medium, first balance the atoms as is done in acidic medium. Then for each H+ ions, add an equal number of OH– ions to both sides of the equation. Where H+ and OH– appear on the same side of the equation, combine these to give H2O.

Ø Electrode Processes : l The redox reactions in which the oxidation and reduction half reactions take place in the same container, are

called the direct redox reactions. e.g. When Zn strip is placed in CuSO4 solution, Zn is oxidized to Zn2+ and CuSO4 is reduced to metallic

copper. Such type of cell is called electrolytic cell. l The redox reactions in which the oxidation and reduction half reactions take place in the separate containers,

are called indirect redox reactions. Such type of cell is called electrochemical cell. For e.g. CuSO4 solution is placed in a beaker with a copper strip or rod and ZnSO4 solution is placed in

separate beaker with a zinc strip or rod. In this experiment, two redox couples are represented as: Zn/Zn2+ and Cu2+/Cu The solutions of two beakers are connected by a salt bridge (a U-tube containing a solution of KCl or NH4NO3).

The Zn and Cu rods are connected by a metallic wire with a provision for an ammeter and switch. This set-up is known as Daniell cell. The two half reactions taking place are :

Zn → Zn2+ + 2e– (Oxidation) at anode Cu2+ + 2e– → Cu (Reduction) at cathode

Overall reaction, Zn + Cu2+ → Zn2+ + Cu Ø The standard electrode potential (EΘ) value for each electrode process is a measure of the relative tendency of the

active species in the process to remain in the oxidized/reduced form. Ø A negative EΘ means that the redox couple is a stronger reducing agent than the H+/H2 couple. Ø A positive EΘ means that the redox couple is a weaker reducing agent than the H+/H2 couple.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 61Know the Terms Ø Redox Couple : It is defined as having together the oxidized and reduced forms of a substance taking part in an

oxidation or reduction half reaction. Ø Electrode : It is an electrical conductor used to make contact with a non-metallic part of a circuit. Ø Electrode Potential : The potential associated with each electrode is called electrode potential. Ø Salt Bridge : It is an inverted U-tube that contains an electrolyte and connects the two half cells in a galvanic cell. Ø Standard Electrode Potential : The potential of an electrode in which the concentration of each species taking

part in the electrode reaction is unity (if any gas is involved in the electrode reaction, its pressure is taken as 1 atmosphere) and the temperature is maintained at 298 K is called the standard electrode potential.

Ø Electrochemical Cell : It is a device which is used to convert chemical energy into electrical energy. Ø Electrochemical Series : The series in which different electrodes are arranged in decreasing order of their standard

reduction potentials (EΘ) is called electrochemical series.

Know the Facts Ø All combustion reactions, which make use of elemental dioxygen, as well as other reactions involving elements

other then dioxygen, are redox reaction . Ø All decomposition reactions are not redox reactions. For example, decomposition of calcium carbonate is not a

redox reaction.

Ca CO s CaO s CO g+ + − + − + −

→ +2 4 2

3

2 2 4 2

2( ) ( ) ( )∆

Ø Metal displacement reactions find many applications in metallurgical processes in which pure metals are obtained from their compounds in ores.

Ø All alkali metals and some alkaline earth metals (Ca, Sr and Ba) which are very good reducing agents, will displace hydrogen from cold water.

e.g. Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

Ø Less active metals such as magnesium and iron react with steam to produce dihydrogen gas.

2Fe(s) + 3H2O(g) ∆→ Fe2O3(s) + 2H2(g) steam Ø Many metals, including those which do not react with cold water, are capable of displacing hydrogen from acids. Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

This reaction is used to prepare dihydrogen gas in the laboratory. Ø Very less active metals, which may occur in the native state such as silver (Ag) and gold (Au) do not react even

with hydrochloric acid. Ø Fluorine is so reactive that it can replace chloride, bromide and iodide ions in solution. Infact, fluorine is so

reactive that it attacks water and displaces the oxygen of water. 2H2O(l) + 2F2(g) → 4HF(aq) + O2(g)

It is for this reason that the displacement reaction of chlorine, bromine and iodine using fluorine are not generally carried out in aqueous solution. On the other hand, chlorine can displace bromide and iodide ions in an aqueous solution.

Cl2(g) + 2KBr(aq) → 2KCl(aq) + Br2(l) Cl2(g) + 2KI(aq) → 2KCl(aq)+ + I2(s) Ø Phosphorus, sulphur and chlorine undergo disproportionation in the alkaline medium as shown below : P4(s) + 3OH–(aq) + 3H2O(l) → PH3(g) + 3H2PO2

–(aq) …(1) S8(s) + 12OH–(aq) → 4S2–(aq) + 2S2O3

2–(aq) + 6H2O(l) …(2) Cl2(g) + 2OH–(aq) → ClO–(aq) + Cl–(aq) + H2O(l) …(3) The third reaction describes the formation of household bleaching agents. The hypochlorite ion (ClO–) formed in

the reaction oxidises the colour bearing stains of the substances to colourless compounds. Ø The reaction that takes place in the case of fluorine is an follows : 2F2(g) + 2OH–(aq) → 2F–(aq) + OF2(g) + H2O(l)

Ø In acid-base systems, titration is used to find the strength of one solution against the other using a pH sensitive indicator.

Ø Functions of Salt bridge in an Electrochemical cell : l It completes the cell circuit by connecting the solutions of two half cells, without allowing them to mix with

each other. l It maintains the electrical neutrality of the solutions in the two half cells due to the flow of ions.

62 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XIØ Representation of an Electrochemical Cell Flow of electrons Flow of current Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Left Salt Right

electrode bridge electrode (Anode) (Cathode) Oxidation Reduction (Negative) (Positive)

Ø A metal having more negative value of EΘ (Standard electrode potential) can displace the metal having less negative value of EΘ from their salt solution.

e.g. Zn can displace Cu from CuSO4 solution (Zn2+/Zn = – 0.76V, Cu2+/Cu + +0.34V) So, Zn is more reactive than Cu. Ø The feasibility of a reaction is decided by its electromotive force (emf) or cell potential. i.e. E°cell = E°cathode – E°anode l A positive value of E°cell suggests that the cell reaction is feasible. l A negative or zero value of E°cell suggests that the cell reaction does not feasible. Ø Highly reactive metals (i.e. having negative value of EΘ) displace hydrogen from acids. For e.g. Metals like Mg, Al, Zn, Fe etc. can replace hydrogen from HCl or H2SO4. However, metals with positive values of EΘ are not capable to replace hydrogen from mineral acids.

qq

Chapter - 9 : Hydrogen

TOPIC-1Dihydrogen

Quick Review Ø Dihydrogen is the most abundant element of this universe. It is estimated about 70 % of the total mass of universe. Ø Hydrogen is the first element in the periodic table and the lighest element with just one proton and one electron. Ø It resembles both alkali metals and halogens and also differs with them in certain respects. Therefore, it is given a

unique position in periodic table. Ø Hydrogen has three isotopes :

(a) Protium, 11H (b) Deutrium, 21H or D

(c) Tritium, 31H or T

Ø Tritium is radioactive and emits low energy b particles. (t½ = 12.33 years). Ø These isotopes show similar chemical proerties as they have same electronic configuration but with little difference

in their rate of reactions due to different bond enthalpies. However, they show considerable difference in their physical properties due to their large mass differences.

Ø Hydrogen can be prepared by various methods : (i) By the action of water on various methods :

2Na + 2H2O (cold) → 2NaOH + H2 Mg + 2H2O (hot) → Mg(OH)2 + H2 3Fe + 4H2O → Fe3O4 + 4H2

(Red hot) (Steam) (ii) By the action of metals on acids and bases. Zn + H2SO4 → ZnSO4 + H2(g) 2Na + 2HCl → 2NaCl + H2(g) 2Al + 2NaOH + 2H2O → 2NaAlO2 + 3H2

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 63 (iii) On large scale, it can be prepared by (a) electrolysis of water (b) from water gas and (c) from hydrocarbons.

(a) 2H O( ) acid/bases

2H (g)+O2 2 2lElectrolysis

Tracesg → ( )

(b) C(s) + H O(g) (g)+ H2 2

water gas or syn gas

→ CO g( )� ��

CO + H2O → CO2(g) + H2(g)

(c) C H + H O1270K

NiCO(g) H2 +2 2 2n n n n n → + +( )2 1

e.g., CH H O(g)K

NiCO(g) H4 2 2( ) ( )g g+ → +

12703

Ø Properties of Dihydrogen— Physical Proerties— Ø Dihydrogen is lighter than air and insoluble in water Ø It is colourless, odourless, tasteless and combustible gas. Ø The H–H bond dissosiation enthalpy is the highest for a single bond between two atoms of any elements. Chemical Properties H2(g) + X2(g) → 2HX(g) (X = F, Cl, Br, I)

2H2(g) + O2(g) → 2H2O(l)

2H2(g) + O2(g) Catalyst or

Heating → 2H2O(l)

3H2(g) + N2(g) 637K, 200 atm

Fe → 2NH3(g)

H2(g) + 2M(g) → 2MH(s) M = alkali metal

H2(g) + Pd2+(aq) → Pd(s) + 2H+(aq)

yH2(g) + MxOy(s) → xM(s) + yH2O(l)

H2 + CO + RCH = CH2 → RCH2CH2CHO

H2 + RCH2CH2CHO → RCH2CH2 CH2OH

Ø Uses of Dihydrogen : It is used for hydrogenation of oils, in the manufacture of metal hydrides and as a rocket fuel in space research. It is used to reduce heavy metal oxide to metals and in oxy-hydrogen torch. It is also used in fuel cells to generate electrical energy. It is also used in the synthesis of ammonia which is used in the manufacture of nitric acid and nitrogenous fertilizers.

Know the Terms Ø Water Gas—The mixture of CO and H2 is called water gas. Ø SynGas—It is a mixture of CO and H2 which is used for synthesis of methanol and a number of hydrocarbons.

It is also called synthesis gas. Ø Coal Gasification—The process of producing ‘Syn gas’ from coal is known as ‘coal gasification’. Ø Water-Gas shift reation—The reaction in which the production of dihydrogen can be increased by reacting CO

of syn gas mixture with steam in the presence of iron chromate as catalyst is called water gas shift reation.

CO(g) H O(g) CO H2

KCatalyst 2+ → +673

2(g) ( )g

(Carbon dioxide is removed by scrubbing with sodium arsenite solution.)

Know the Facts Ø Hydrogen has the simplest atomic structure among all the elements around us in nature. Ø Its atomic form consists of only one proton and one electron. Ø In elemental form, it exists as a diatomic (H2) molecule and is called dihydrogen. Ø The electronic configuration of hydrogen is 1s1. Ø Loss of electron from hydrogen atom results in nucleus (H+) of ∼1.5 × 10–3 pm size. This is extremely small as

compared to normal atomic and ionic sizes of 50 to 200 pm. As a consequence, H+ does not exist freely and is always associated with other atoms or molecules. So, it is unique in behaviour and is, therefore, best placed separately in the periodic table.

Ø Dihydrogen is the most abundent element in the universe (70%of the total mass of the universe).

64 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø It is the principal element in the solar atmosphere. Ø The giant planets Jupiter and Saturn consist mostly of hydrogen. Ø Due to light nature of hydrogen, it is much less abundant (0.15 % by mass) in the earth’s atmosphere. However.

in the combined form, it constitutes 15.4% of the earth’s crust and the oceans. Ø An American scientist, Harold C. Urey (1934) got Nobel Prize for separating isotope of hydrogen having mass

number 2 by physical methods.

Ø The predominant form of isotope of hydrogen is protium 11H( ).

Ø Terrestrial hydrogen contains 0.0156 % of deuterium 12 H( ) mostly in the form of heavy water (D2O).

Ø Presently ~ 77% of the industrial dihydrogen is produced from petrochemicals, 18% from coal, 4% from electrolysis of a queous solutions and 1% from other sources.

Ø The H – H bond dissociation enthalpy is the highest for a single bond between two atoms of any element. Ø It is relatively inert at room temparature due to high H – H bond enthalpy. Ø The atomic hydrogen can be produced at a high temperature in an electric arc or under UV radiations. It can

combine with almost all elements. Ø Hydrogen energy does not produce any pollution and releases greater energy per unit mass of fuel in comparison

to gasoline and other fuels.

TOPIC-2Hydrides and Water

Quick Review Ø Hydrides : Dihydrogen under certain reaction conditions form hydrides. There are three types of hydrides : (i) Ionic hydrides are the hydrides of s-block element. e.g., LiH, NaH, BeH2, MgH2. (ii) Covalent and molecular hydrides are molecular compounds with p-block elements. e.g., CH4, NH3, H2O and

HF. These can be further classified as : (a) Electron : deficient hydrides : e.g., Diborane (B2H6). (b) Electron : precise hydrides : e.g., Methane (CH4) and other hydrides of group-14 element. (c) Electron : rich hydrides : Hydrides of group 15-17 elements form these types of hydrides. e.g., NH3,

H2O, HF etc. (iii) Metallic or non-stoichiometric hydrides are hydrides of many d-block and f-block elements. Metals of group

6 (except ZrH), 7, 8 and 9 do not form hydrides. e.g., LaH2.87, YbH2.55, TiH1.5-1.8, ZrH1.3-1.75, VH0.56, NiH0.6-0.7, PdH0.6-0.8 etc.

ØWater is the most common and abundant substance that is of great chemical and biological significance. Ø Water has a high specific heat, thermal conductivity, surface tension, dipole moment and dielectric constant etc. Ø Ice (solid) has open cage type structure because of that its density is less than liquid water. Therefore, ice floats on

liquid water. Ø In the gas phase, water is a bent molecule with a bond angle of 104.5° and O-H bond length of 95.7 pm. Its

structure is represented as below:

Ø In liquid phase, water molecules are associated together by hydrogen bonds. Ø The crystalline form of water is ice. Ø Chemical Properties of Water : (i) Amphoteric Nature : Water acts as an acid and base. H2O (l) + NH3 (aq) � OH– (aq) + NH4

+ (aq) acid base H2O(l) + H2S(aq) � H3O+(aq) + HS–(aq) base acid (ii) Reduction Reaction : 2H2O + 2Na → 2NaOH + H2

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 65 (iii) Hydrolysis Reaction : SiCl4 + 2H2O → SiO2 + 4HCl

P4O10(s) + 6H2O(l) → 4H3PO4(aq)

N3–(s) + 3H2O(l) → NH3(g) + 3OH–(aq)

(iv) Water can be reduced to hydrogen by highly electropositive metals.

(v) Water is oxidised to O2 during photosynthesis and with fluorine.

6CO2(g) + 12H2O(l) → C6H12O6(aq) + 6H2O(l) + 6H2O(g)

2F2(g) + 2H2O(l) → 4H+(aq) + 4F–(aq) + O2(g)

(vi) Hydrates formation : Many salts can be crystallised as hydrates salt from their aqueous solution. Coordinated water e.g, [Cr(H2O)6]3+ 3CI–

Interstitial water, e.g., BaCI2.2H2O Hydrogen bonded water i.e., [Cr(H2O)4]2+ SO4

2– . H2O in CusSO4.5H2O. Ø Hard and soft water : Hard water does not give lather with soap. Soft water easily forms lather with soap. Hard water forms scum or precipitate with soap. C17H35COONa + M2+ → (C17H35COO)2 M ↓ + 2Na+

(Soap) From hard water Scum Sodium steareate M is Ca or Mg.

Ø Disadvantages of Hard Water : (i) It is unsuitable for laundary. (ii) It is harmful for boilers. (iii) It reduces the efficiency of boiler. Ø There are two types of Hardness : (i) Temporary Hardness (ii) Permanent Hardness (i) (a) Temporary hardness of water is due to the presence of bicarbonates of magnesium and calcium. (b) Temporary hardness of water can be removed by boiling during which soluble M(HCO3)2 is converted

to insoluble Mg(OH)2 and Ca(HCO3)2 is changed to insoluble CaCO3. It can also be removed by Clark’s method in which calculated amount of lime is added in water that precipitates calcium carbonate and magnesium hydroxide.

(ii) (a) Permanent hardness of water is due to the presence of chlorides and sulphates of magnesium and calcium.

(b) Permanent hardness of water can be removed by treatment with washing soda (sodium carbonate), or by Calgon’s method (addition of sodium hexametaphosphate) or by ion-exchange method or by synthetic resin method.

Know the Terms Ø Hydrides—The binary compounds of hydrogen and other elements (except noble gases) are called hydrides. Ø Ionic Hydrides—The stoichiometric compounds of dihydrogen and highly electropositive elements are called

ionic hydrides. Ø Hard Water—Water containing soluble salts of calcium and magnesium is called hard water. Ø Soft Water—Water free from soluble salts of calcium and magnesium is called soft water. Ø Hardness—The capacity of water which restricts the formation of lather with soap is called hardness.

Know the Facts Ø Significant covalent character is found in lighter metal hydrides such as LiH, BeH2 and MgH2. Ø BeH2 and MgH2 are polymeric in nature. Ø The ionic hydrides are crystalline, non-volatile and non-conducting in solid state. Ø Ionic hydrides in molten state conduct electricity and on electrolysis liberate dihydrogen gas at anode, which can

firms the existence of H– ion.

2H– (melt) anode → H2 (g) + 2e–

Ø Saline hydrides react violently with water to produce dihydrogen gas.

NaH(s) + H2O(aq) → NaOH(aq) + H2(g)

Ø Lithium hydride is rather unreactive at moderate temperature with oxygen and chlorine. So, it is used to synthesize other useful hydrides.

e.g. 2LiH + B2H6 → 2LiBH4

8LiH + Al2Cl6 → 2LiAlH4 + 6LiCl

66 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø Electron-precise compounds have the required number of electrons to write their conventional lewis structures.

All elements of group 14 form such compounds (e.g., CH4) which are tetrahedral in geometry. Ø Electron-rich hydrides have excess electrons which are present as lone pairs. Ø The property of absorption of hydrogen on transition metals is widely used in catalytic reduction or hydrogenation

reactions for the formation of large number of compounds. Ø Some of the metals (e.g. Pd, Pt) can accommodate a very large volume of hydrogen and so can be used as its

storage media. This property has high potential for hydrogen storage and as a source of energy. Ø Human body has about 65% and some plants have as much as 95 % water. Ø Water is a colourless and tasteless liquid. Ø The high heat of vaporisaton and heat capacity of water are responsible for moderation of the climate and body

temperature of living beings. Ø Water is an excellent solvent for transportation of ions and molecules required for plant and animal metabolism. Ø Due to hydrogen bonding with polar molecules, even covalent compounds like alcohol and carbohydrates

dissolve in water. Ø At atmospheric pressure, ice crystallises in the hexagonal form, but at very low temperatures, it condenses to

cubic form. Ø Density of ice is less than that of water. So, ice cube floats on water. It is due to presence of inter-molecular

hydrogen bonding in water molecules. Ø In winter season, ice formed on the surface of a lake provide thermal insulation which ensures the survival of the

aquatic life. This fact is of great ecological significance. Ø Hydrogen bonding gives ice a rather open type structure with wide holes. These holes can hold some other

molecules of appropriate size interstitially. Ø Sodium hexametaphosphate (Na6P6O18) is commercially called calgon. It is used in the internal treatment of hard water. Ø Hydrated sodium alumino silicate [Na2O.Al2O3.xSiO2. yH2O] is called zeolite or pemutit. It is used in the treatment

of hard water. Ø The unit of hardness is ppm (parts per million) or mg/L.

TOPIC-3Hydrogen Peroxide, Heavy Water and Dihydrogen as a Fuel

Quick Review Ø H2O2 (Hydrogen Peroxide) : (i) Merck’s Process (Preparation)

Na2O2 + H2SO4 —→ Na2SO4 + H2O2

(ii) Industrially, it is manufactured by the auto-oxidation of 2-alkylanthraquinols.

2-ethylanthraquinol

O

H Pd2

2

(air)

/� ⇀��↽ ��

H2O2 + (Oxidised product)

Ø Chemical Properties : (i) 2Fe2+ + 2H+ + H2O2 → 2Fe3+ + 2H2O [Oxidising Action in Acidic

Medium] (ii) Reducing action in acidic medium 2MnO4

– + 6H+ + 5H2O2 → 2Mn2+ + 8H2O + 5O2

(iii) Oxidising action in basic medium 2Fe2+ + H2O2 → 2Fe3+ + 2OH–

(iv) Reducing action in basic medium 2MnO4

– + 3H2O2 → 2MnO2 + 3O2 + 2H2O + 2OH–

Ø Storage of H2O2 : H2O2 decomposes slowly on exposure to light. Hence it is stored in wax lined glass or plastic bottles in dark.

2H2O2 (l) → 2H2O (l) + O2 (g)

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 67 Ø Structure of H2O2 : (i) H2O2 structure (gas phase) dihedral angle 111.5°, (ii) H2O2 (solid phase at 110 K. The dihedral angle is reduced to 90.2°).

Ø Uses of Hydrogen Peroxide : (i) It is used in the industries as a bleaching agent. (ii) It is used as a hair bleach and mild disinfectant.

Ø Heavy water (D2O) : Heavy water is the form of water consisting of isotope of hydrogen (deuterium). It is used as moderator in nuclear reactor.

Ø Volume strength of H2O2 : A 30% solution of H2O2 is marketed as ‘‘100 volume’’ hydrogen peroxide and one millilitre of 30% H2O2 solution will give 100 ml of oxygen at STP.

Ø Dihydrogen as a fuel : Dihydrogen releases large quantities of heat on combustion. Ø Hydrogen economy is an alternative source of energy. The basic principle of hydrogen economy is the

transportation and storage of energy in the form of liquid or gaseous dihydrogen. Ø Advantage of hydrogen economy is that energy is transmitted in the form of dihydrogen and not as electric

power.

Know the Terms Ø Fuel—Any combustible substance which produces heat on its complete combustion is called fuel. Ø 10 Volume solution of H2O2—It means that 1 L of H2O2 solution will give 10 L of oxygen at STP.

Know the Facts Ø In the pure state, H2O2 is an almost colourless (very pale blue) liquid. Ø H2O2 is miscible with water in all proportions and forms a hydrate H2O2.H2O (m.p. 221 K). Ø A 30 % solution of H2O2 is marketed as ‘100 volume’ H2O2. It means that 1 mL of 30% H2O2 solution will give

100mL of oxygen at STP. Ø Commercially, marketed sample in 10 V, which means that the sample contains 3 % H2O2. Ø H2O2 acts as an oxidising as well as reducing agent in both acidic and alkaline media. Ø Urea can be added in H2O2 as a stabilizer. Ø H2O2 is kept away from dust because dust can induce explosive decomposition of the compound. Ø H2O2 is used as an antiseptic which is sold in the market as perhydrol. Ø H2O2 is used in Environmental (Green) Chemistry. For example, in pollution control treatment of domestic and

industrial effluents, oxidation of cyanides, restoration of aerobic conditions to sewage wastes etc. Ø Heavy water (D2O) can be prepared by exhaustive electrolysis of water or as a by-product in some fertilizer

industries. Ø On mass basis, dihydrogen can release more energy than petrol (about 3 times). Ø Pollutants in combustion of dihydrogen will be less than petrol. The only pollutants will be oxides of dinitrogen

(due to the presence of dinitrogen as impurity with dihydrogen). This can be minimized by injecting a small amount of water into the cylinder to lower the temperature so that the reaction between dinitrogen and dioxygen may not take place.

Ø It is for the first time in the history of India that a pilot project using dihydrogen as a fuel was launched in October 2005 for running automobiles.

Ø Initially 5 % dihydrogen has been mixed in CNG for use in four-wheeler vehicles. Ø Now a days, dihydrogen is used in fuel cells for the generation of electric power.

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Chapter - 10 : s-Block Elements (Alkali and Alkaline Earth Metals)

TOPIC-1 Alkali Metals and Their Compounds

Quick Review Ø Alkali metals : The group-1 elements are lithium, sodium, potassium, rubidium, caesium and francium. Ø The general electronic configuration of alkali metals is ns1. Ø Lithium is the lightest metal. Ø Lithium shows diagonal relationship with Mg. Ø Sodium is the most abundant alkali metal in the earth crust. Ø Alkali metals have lowest ionization enthalpies and largest atomic radii in their respective periods. Ø The hydration enthalpies of alkali metal ions are in the order :

Li+ > Na+ > K+ > Rb+ > Cs+.

Ø Alkali metals behave as strong reducing agents due to their low ionization enthalpies. Ø All alkali metals impart characteristic colour to the flame. Li—caramine red, Na—golden yellow, K—pale violet,

Rb—reddish violet and Cs–sky blue. Ø Caesium, due its very low ionization enthalpy, is used in photoelectric cells. Ø Alkali metals, due to their low melting and boiling points have very weak metallic bond. This is why they can be

cut with the help of a knife. Ø The reactivity of alkali metals increases from lithium to caesium. Ø Alkali metal hydrides are ionic in nature and are strong reducing agents. Ø Lithium forms monoxide Li2O even with excess of oxygen. Na forms peroxide while K, Rb and Cs form superoxides

on reaction with excess oxygen. Ø The basic character of alkali metal hydoxides increases on descending the group. Ø Alkali metals dissolve in liquid ammonia giving deep blue solutions which contain ammoniated cations and

ammoniated electrons. M+ (x+y)NH M(NH ) e NHx y3 3 3→ ++ −[ ] [ ( ) ] (where M = alkali metal) Ø The blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region of

light and imparts blue colour to the solution. Ø Lithium exhibits anomalous behaviour due to its small size, high ionization enthalpy and greater polarizing

power. Ø Some Important Compounds of Sodium– l Sodium Carbonate : Ø Its chemical formula is Na2CO3 . 10H2O . It is also known as washing soda. Ø It is generally prepared by Solvay process. Ø The following reactions are involved in the preparation of sodium carbonate–

22

3 2 4 2 3

4 2 3 3

3

NH H O + CO NH CONH CO H O+CO NH HCO

NH HCO Na

2

2 2 4

4

+ →+ →+

( )( )

CCl NH Cl+NaHCOSodium bicarbonate→ 4 3

Ø Sodium bicarbonate crystals separate which are heated to give sodium carbonate. 2NaHCO Na CO CO H O3 2 3 2 2→ + + Ø In this process, NH3 is recovered when the solution containing NH4Cl is treated with Ca(OH)2. Calcium Chloride

is obtained as a by-product. 2 24 2NH Cl Ca(OH) NH CaCl H O3 2 2+ → + + Ø Solvay process cannot be used for the manufacture of potassium carbonate because potassium hydrogen

carbonate (KHCO3) is too soluble to be precipitated by the addition of ammonium hydrogen carbonate to a saturated solution of potassium chloride.

Ø Sodium carbonate is a white crystalline solid which exists as a decahydrate, Na2CO3 . 10 H2O.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 69 Ø Carbonate part of sodium carbonate gets hydrolysed by water to form an alkaline solution.

CO H O HCO OH2 332− − −+ → +

Ø Sodium carbonate is used in water softening, laundering and cleaning. It is also used in the manufacture of glass, soap, borax and caustic soda.

l Sodium Chloride : Ø Its chemical formula is NaCl. Ø The most abundant source of sodium chloride is sea water which contains 2.7 to 2.9% by mass of the salt. Ø Crude sodium chloride is obtained by crystallisation of brine solution. It contains CaCl2, Na2SO4, CaSO4 and

MgCl2 as impurties. Ø To obtain pure NaCl, the crude salt is dissolved in minimum amount of water and filtered to remove insoluble

impurities. The solution is then saturated with HCl gas. Crystals of pure NaCl separates out. CaCl2 and MgCl2, being more soluble than NaCl, remain in solution.

Ø Sodium chloride melts at 1081 K. Ø Sodium chloride is used as a common salt for domestic purpose. It is also used for the preparation of Na2O2,

NaOH and Na2CO3. l Sodium Hydroxide : Ø Its chemical formula is NaOH. It is commonly known as caustic soda. Ø It is generally prepared commercially by the electrolysis of sodium chloride in Castner – Kellner cell. The reactions

take place are as follow – At athode: Na e Na-Hg

Sodium amalgam

+ Hg+ →−

At anode: Cl Cl e-−→ +12 2

Ø The sodium – amalgam is treated with water to give sodium hydroxide and hydrogen gas.

2 2 2Na-Hg+2H O 2NaOH+2Hg+H→

Ø It is a white, translucent solid. Ø Its melting point is 591 K. Ø It is readily soluble in water to form a strong alkaline solution. Ø Crystals of NaOH are deliquescent. Ø The sodium hydroxide solution at the surface reacts with CO2 in the atmosphere to form Na2CO3. Ø It is used in the manufacture of soap, paper, artificial silk and a number of chemicals. Ø It is also used as laboratory reagent. l Sodium Hydrogen Carbonate : Ø Its chemical formula is NaHCO3. It is commonly known as baking soda. It is due to the fact that it decomposes on

heating to generate bubbles of CO2 (leaving holes in cakes or pastries and making them light and fluffy). Ø It is formed by saturating a solution of sodium carbonate with CO2. The white crystalline powder of sodium

hydrogen carbonate, being less soluble, gets separated out. Na CO H O+CO NaHCO2 3 2+ →2 32

Ø It is a mild antiseptic for skin infections. Ø It is also used in fire extinguishers. l Biological Importance of Sodium and Potassium : Ø Sodium ions are found primarily on the outside of cells, in blood plasma and in the interstitial fluid which

surrounds the cells. Ø Sodium ions participate in the transmission of nerve signals, in regulating the flow of water across cell membranes

and in the transport of sugars and amino acids into cells. Ø Sodium and potassium are similar chemically but they differ quantitatively in their ability to penetrate cell

membranes, in their transport mechanisms and in their efficiency to activate enzymes. Ø Potassium ions are the most abundant cations in cell fluids, where they activate many enzymes, participate in

the oxidation of glucose to produce ATP (Adenosine triphosphate) and, with sodium, are responsible for the transmission of nerve signals.

Ø There is a considerable variation in the concentration of sodium and potassium ions found on the opposite sides of cell membrane.

e.g. In blood plasma, sodium is present to the extent of 143 m mol L–1 while potassium is only 5 m mol L–1. On the other hand, in RBCs these concentrations vary to 10 m mol L–1 (Na+) and 105 m mol L–1 (K+).

Ø Sodium – potassium pump operates across the cell membranes which consumes more than one–third of the ATP used by a resting animal and about 15 kg per 24 hrs in a resting human.


Know the Terms Ø s–block elements : The elements in which the last electron enters the outermost s–orbital are called s–block

elements. Ø Alkali metals : Group 1 elements of the periodic table are collectively known as alkali metals. These are so called

because they form hydroxides on reaction with water which are strongly alkaline in nature. Ø Diagonal Relationship : It is said to exist between certain pairs of diagonally adjacent elements in the second and

third periods of the periodic table. Ø Hydration Enthalpy : The amount of energy released when one mole of gaseous ions combine with water to form

hydrated ions is known as hydration enthalpy.

Know the Facts Ø Among the alkali metals, sodium and potassium are abundant and Li, Rb and Cs are much less abundant. Ø Francium (Fr) is highly radioactive, its longest – lived isotope 223Fr has a half–life of only 21 minutes. Ø All the alkali metals are silvery white, soft and light metals. Ø Due to large size, alkali metals have low density which increases down the group from Li to Cs. However,

potassium is lighter than sodium. Ø The alkali metals are highly reactive due to their large size and low ionisation enthalpy. Ø The alkali metals tarnish in dry air due to formation of their oxides which in turn react with moisture to form

hydroxides.

4 2

2

2 2

2

Li+O Li OLithium oxide

Na+O Na OSodium peroxide

M+O M

2 2

2

→

→

→ OOSuperoxide

M K, Rb, Cs

2

( )=

Ø In all oxides, the oxidation state of alkali metal is +1. Ø The alkali metals react with water to form hydroxide and dihydrogen.

2 2 22M+2H O M OH H2→ + ++ −

(M= alkali metal) Ø The alkali metals react with dihydrogen at about 673 K (Li at 1073 K) to form hydrides.

2 22M+H M H→ + −

(M= alkali metal) Ø The alkali metals react vigorously with halogens to form ionic halides (M+X–) Ø The alkali metals are strong reducing agents. Lithium is most powerful reducing agent. Ø Lithium is used to make useful alloys such as white metal with lead. Ø It is used in thermonuclear reactions. Ø It is used to make electrochemical cells. Ø Liquid sodium metal is used as a coolant in fast breeder nuclear reactors. Ø Potassium has a vital role in biological systems. Ø KCl is used as a fertilizer. Ø KOH is used in the manufacture of soft soap. Ø Caesium is used in devising photoelectric cells. Ø Relative order of melting and boiling points among halides of alkali metals : fluoride > chloride > bromide > iodide Ø The low solubility of LiF in water is due to its high lattice enthalpy while the low solubility of CsI is due to smaller

hydration enthalpy of its two ions. Ø The m.p. and b.p. of lithium are higher than other alkali metals. Ø LiCl is deliquescent and crystallises as a hydrate, LiCl . 2H2O while other alkali metal chlorides do not form

hydrates. Ø Both Li and Mg are harder and lighter than othr elements in the respective groups. Ø Sodium chloride has a solubility of 36 g in 100 g of water at 273 K. The solubility does not increase appreciably

with increase in temperature. Ø A typical 70 kg man contains about 90 g of sodium and 170 g of potassium compared with only 5g of iron and

0.06 g of copper. Ø Due to large negative electrode potentials, alkali metals are stronger reducing agent than alkaline earth metals.


TOPIC-2 Alkaline Earth Metals and Their Compounds

Quick Review Ø Alkaline earth metals correspond to Group 2 which comprise Be, Mg, Ca, Sr, Ba and Ra. Ø The first element beryllium (Be) differs from the rest of the members and shows diagonal relationship to

aluminium. Ø The general electronic configuration of alkaline earth metals is ns2. Ø Like alkali metals, the compounds of these elements are also predominantly ionic. Ø The atomic and ionic radii of the alkaline earth metals are smaller than those of the corresponding alkali metals

in the same periods. This is due to the increased nuclear charge in these elements. Ø With in the group, the atomic and ionic radii increase with increase in atomic number. Ø The alkaline earth metals have low ionisation enthalpies due to fairly large size of the atoms. The ionisation

enthalpy decreases down the group due to increase in atomic size. Ø The first ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1

metals which is due to their small size as compared to the corresponding alkali metals. Ø Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size

down the group.

Be2+ > Mg2+ > Ca2+ > Sr2+ > Ba2+

Ø The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions. Ø The alkaline earth metals are silvery white, lustrous and relatively soft but harder than the alkali metals. Ø The melting and boiling points of alkaline earth metals are higher than the corresponding alkali metals due to

smaller sizes. Ø The alkaline earth metals are less reactive than the alkali metals. Ø Calcium, strontium and barium are readily attacked by air to form the oxide and nitride. Ø All the alkaline earth metals combine with halogen at elevated temperatures forming their halides. Ø All the alkaline earth metals except Be combine with hydrogen upon heating to form their hydrides, MH2. Ø The alkaline earth metals readily react with acids liberating dihydrogen.

M+2HCl MCl H2 2→ +

Ø Like alkali metals, the alkaline earth metals are strong reducing agents. However, their reducing power is less than those of their corresponding alkali metals.

Ø Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.

M NH M(NH ) e(NH )3+ + → ++ −( ) [ ] [ ]x y x y32

32

Ø Metallic beryllium is used for making windows of X-ray tubes. Ø A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine. Ø MgCO3 is an ingredient of tooth paste. Ø Radium salts are used in radiotherapy i.e. in the treatment of cancer. Ø The alkaline earth metals burn in oxygen to form the monoxide (MO). Ø BeO is amphoteric while oxides of other elements are ionic in nature. Ø All these oxides except BeO are basic in nature and react with water to form sparingly soluble hydroxides. MO+H O M(OH)2 → 2

where M = alkaline earth metals Ø All the halides of alkaline earth metals except beryllium halides are ionic in nature. Beryllium halides are covalent

in nature. Ø The alkaline earth metals form salts of oxoacids like carbonates, sulphates, nitrates etc. Ø The anomalous behaviour of beryllium is due to its small atomic and ionic size.

Some Important Compounds of Calcium - l Calcium Oxide : Ø Its chemical formula is CaO. It is commonly known as quick lime. Ø It is prepared commercially by heating lime stone in a rotary kiln at 1070–1270 K.

CaCO CaO+CO3 2∆� ⇀��↽ ��

Ø It is a white amorphous solid. Ø Its melting point is 2870 K.

72 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Ø On exposure to atmosphere, it absorbs moisture and carbon dioxide.

CaO H O Ca OHCaO CO CaCO

2+ →+ →

( )2

2 3 Ø It is used as a raw material for manufacturing cement and is the cheapest form of alkali. Ø It is also used in the manufacture of dye stuffs. l Calcium Hydroxide : Ø Its chemical formula is Ca(OH)2. It is commonly known as slaked lime. Ø It is prepared by adding water to quick lime (CaO). Ø It is a white amorphous powder. Ø The aqueous solution of Ca(OH)2 is known as lime water and a suspension of slaked lime in water is called milk

of lime. Ø When CO2 is passed through lime water, it turns milky due to the formation of calcium carbonate.

Ca OH CO CaCO H O( )2 2 3 2+ → + Ø On passing excess of CO2, the precipitate dissolves to form calcium hydrogen carbonate. CaCO CO H O Ca HCO3 2 2 3 2+ + → ( ) Ø Milk of lime reacts with chlorine to form hypochlorite, a constituent of bleaching powder.

2 2 22 2 2 2 2Ca OH Cl CaCl Ca OCl H O

( ) ( )+ → + + Bleaching powder

It is used in the preparation of mortar, a building material. Ø It is also used in making glass, in tanning industry etc l Calcium Carbonate : Ø Its chemical formula is CaCO3 . It occurs in nature in the form of lime stone, chalk, marble etc. Ø It can be prepared by passing CO2 through slaked lime. Ca OH CO CaCO H O( )2 2 3 2+ → + Ø It is also prepared by addition of sodium carbonate to calcium chloride. CaCl2 + Na2CO3 → CaCO3 + 2 NaCl Ø It is a white fluffy powder. Ø On heating calcium carbonate at 1200 K, it decomposes to evolve carbon dioxide. CaCO CaO COK

31200

2 → + Ø It is used as a building material in the form of marble and also in the manufacture of quick lime. Biological Importance of Mg and Ca- l All enzymes that utilise ATP in phosphate transfer require magnesium as the cofactor. l The main pigment for the absorption of light in plants is chlorophyll which has magnesium. l About 99% of body calcium is found in bones and teeth. l Calcium plays an important role in neuromuscular function, interneuronal transmission, cell membrane in-

tegrity and blood coagulation.

Know the Terms Ø Milk of Magnesia : A suspension of magnesium hydroxide in water is called milk of magnesia. Ø Slaking of Lime : The process in which the addition of limited amount of water breaks the lump of lime is called

slaking of lime.

Know the Facts Ø The flame test for Ca, Sr and Ba is helpful in their detection in qualitative analysis and estimation by flame pho-

tometry. Ø Calcium and barium have been used to remove air from vacuum tubes due to their reactivity with oxygen and

nitrogen at high temperatures. Ø Beryllium hydroxide is amphoteric in nature as it reacts with acid and alkali both.

Be(OH) OH Be(OH)

Beryllate ion2 4

22+ →− −[ ]

Be(OH) HCl+2H O [Be(H O) ]Cl22 2 4 22+ →

Ø Beryllium chloride has a chain structure in the solid state as shown below -

Be

Cl

Cl

Be

Cl

Cl

Be

Cl

Cl

Ø In the vapour phase, BeCl2 tends to form a chloro bridged dimer which dissociates into the linear monomer at high temperatures of the order of 1200 K.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 73 Ø The fluorides of alkaline earth metals are relatively less soluble than the corresponding chlorides due to their high

lattice energies. Ø The greater hydration enthalpies of Be2+ and Mg2+ ions overcome the lattice enthalpy factor and so their

sulphates are soluble in water. Ø Beryllium and aluminium ions have strong tendency to form complexes i.e. BeF4

2–, AlF63–

Ø An adult body contains about 25 g of Mg and 1200 g of Ca compared with only 5 g of iron and 0.06 g of copper. Ø CaSO4 . 1/2H2O is commonly known as Plaster of Paris. Ø CaSO4 . 2H2O is commonly known as gypsum. Ø Cement is a mixture of lime and clay having silica along with oxides of Al, Fe and Mg. Ø The thermal stability of carbonates of alkaline earth metals increases with increase in size of cation.

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Chapter - 11 : p-Block Elements

TOPIC-1General Introduction to p-block Elements and Group 13 Elements and Important Compounds of Boron

Quick ReviewØ Group 13 to 18 of the periodic table of elements constitute the p–block. The p–block contains metals, metalloids

as well as non-metals. Ø Electronic configuration—They have general valence shell electronic configuration ns2 np1–6.

Ø GROUP 13 ELEMENTS—The boron family: Boron is a typical non-metal, aluminium is a metal but shows many chemical similarities to boron, gallium, indium and thallium are almost exclusively metallic in character.Ø Electronic configuration—The outer electronic configuration of these elements is ns2 np1.Ø Occurrence-Boron is a fairly rare element but aluminium is the most abundant metal and the third most abun-

dant element in the earth’s crust.Ø Atomic properties:

(i) Atomic radii—Atomic radius increases down the group with thallium (Tl) having the largest atomic radius. (Atomic radius of gallium is less than aluminium due to poor shielding effect of d-orbital electrons.)

(ii) Ionisation enthalpy—First ionisation energy decreases down the group. Some discontinuity in the ionisation enthalpy values observed between Al and Ga, and between In and Tl are due to inability of d- and f-electrons, which have low screening effect, to compensate the increase in nuclear charge.

(iii) Electronegativity—Down the group, electronegativity first decreases from B to Al and then increases marginally. This discrepancy is due to aberrant trends in atomic size.

Ø Physical properties—Boron is non-metallic and hard solid. It exists in many allotropic forms. Due to very strong crystalline lattice, it has high melting point. Rest of the members are soft metals with low melting point and high electrical conductivity.Ø Oxidation state—B and Al show an oxidation state of +3 only while Gallium, Indium and Thallium show oxida-

tion states of both +1 and +3. Further due to inert pair effect, as we move down the group, the stability of +3 oxidation state decreases while that of +1 oxidation state increases.Ø These are electron deficient elements, hence act as good Lewis acids. The tendency to behave as Lewis acids

decreases down the group.Ø Chemical Reactivity:

(i) Reactivity towards Air: Boron and Aluminium react at high temperature with oxygen and nitrogen to form oxides and nitrides respectively. Oxide of boron is acidic, oxides of aluminium and gallium are amphoteric, while oxides of indium and thallium are basic in nature.

2E(s) + 3 O2 (g) ∆ → 2E2O3 (s),

2E(s) + N2 (g) ∆ → 2EN (s), (here E=element) (ii) Reactivity towards Acids and Alkalies: Boron is non-reactive at moderate temperature to acids and alkalies,

whereas aluminium reacts with both acids and bases. Thus it shows amphoteric character. 2Al(s) + 6HCl(aq) → 2Al3+ (aq) + 6Cl–(aq)+ 3H2(g) 2Al(s) + 2NaOH (aq) + 6H2O (l) → 2Na+ [Al (OH) 4] –(aq) + 3H2(g) Sodium tetrahydroxoaluminate (III)

74 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI (iii) Reactivity towards Halogens: These elements react with halogen to form trihalides (except TlI3). 2E(s) + 3X2 (g) → 2EX3 (s) (X = F, Cl, Br, I)

Ø Anomalous Properties of Boron—The first member of group 13, boron differs in many respects from the other members of their respective groups because of small size, high electronegativity and absence of d–orbitals. Due to these factors, certain important trends can be observed in the chemical behaviour of group 13 elements.Ø Some Important Compounds of Boron:

(i) Borax: Borax, Na2B4O7.10H2O is represented as Na2[B4O5(OH)4].8H2O. Aqueous solution of borax is alkaline which is represented as:

Na2B4O7 + 7H2O → 2NaOH + 4H3BO3 Orthoboric acid Properties : On heating, borax first loses water molecules and swells up. On further heating, it turns into a

transparent liquid, which solidifies into glass like material known as borax bead. Na2B4O7.10H2O ∆ → Na2B4O7 ∆ → 2NaBO2 + B2O3 Sodium metaborate Boric anhydride (ii) Orthoboric acid: Preparation : It is prepared by acidifying an aqueous solution of borax. Na2B4O7 + 2HCl + 5H2O → 2NaCl + 4B(OH)3 Boric acid is a weak monobasic acid. It is not a protonic acid but acts as a Lewis acid. B(OH)3 + 2HOH → [B(OH)4]– + H3O+

Properties : On heating, orthoboric acid above 370 K forms metaboric acid, HBO2 which on further heating yields boric oxide, B2O3.

H3BO3 ∆ → HBO2 ∆ → B2O3 Metaboric acid Boric oxide

It has a layer structure in which planar BO3 units are joined by hydrogen bonds.

(iii) Diborane : The simplest boron hydride known is diborane. Preparation : It is prepared as: (i) 4BF3 + 3LiAlH4 → 2B2H6 + 3LiF + 3AlF3

(ii) 2NaBH4 + I2 → B2H6 + 2NaI + H2 (laboratory method) (iii) 2BF3 + 6NaH → B2H6 + 6NaF (Industrial method) Properties: Diborane is a colourless, highly toxic gas that catches fire immediately on exposure to air. Other reactions of

diborane are: (i) B2H6(g) + 6H2O (l) → 2B(OH)3 (aq) + 6H2(g) Boric acid (ii) B2H6 + 2NMe3 → 2BH3.NMe3

B2H6 + 2CO → 2BH3.CO

(iii) 3B2H6 + 6NH3 → 3[BH2 (NH3)2] + [BH4]– ∆ → 2B3N3H6 + 12H2 Borazine Borazine, B3N3H6 is known as inorganic benzene as it is having ring structure with alternate BH and NH groups. In diborane, each boron atom is bonded with two terminal hydrogen atoms. Two other hydrogen atoms form be-

tween two boron atoms and are known as bridging hydrogen atoms. These bonds are three centre—two electron bonds.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 75Ø Uses of Boron, Aluminium and their compounds: (i) Boron fibres are used in making bullet-proof vest and light

material for aircrafts. (ii) Boron-10 isotope is used in nuclear industry. (iii) Boron and boric acid are used in the manufacture of heat resistant glasses, fibre glass and glass-wool. (iv) Aluminium and its alloys are used in tubes, rods, wires, plates or foil etc.

Know the TermsØ p-Block Elements—The elements in which last electron enters the outermost p-orbital are called p-block elements.Ø Inert pair effect—The inert pair effect is the tendency of the two electrons in the outermost atomic s orbital to

remain unionized or unshared in compounds of post-transition metals.

Know the FactsØ Banana Bond: In diborane, three centre—two electron bond occurs which are called bent or banana bond.Ø Inert pair effect increases down a group and thus the elements present in the lower part of the group show lower

oxidation states which is two units less than the highest group oxidation state.Ø Boron shows diagonal relationship with silicon; both are semiconductors metalloids and forms covalent

compounds.Ø Aluminum halides, e.g., AlCl3 is dimer, an important catalyst in organic chemistry have an incomplete octet, acts

as Lewic acid by accepting lone pairs from Lewic bases, forming adduct.Ø Borazine has a cyclic structure similar to benzene and thus is called inorganic benzene.

TOPIC-2Group 14 Elements and Important Compounds of Carbon and Silicon

Quick ReviewØ GROUP 14 ELEMENTS : The Carbon family : Carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb) are

the members of group 14. Carbon is a non-metal, silicon and germanium are metalloids, tin and lead are metals.Ø Electronic configuration : They contain four electrons in the outer most shell i.e. a filled s-orbital and two elec-

trons in the p-orbital. (ns2np2).Ø Occurrence—Carbon is the seventeenth most abundant element by mass in the earth’s crust. Silicon is a very

important component of ceramics, glass and cement. Germanium exists only in traces. Tin occurs mainly as cas-siterite, SnO2 and lead as galena, PbS.Ø Atomic properties :

(i) Atomic radii : There is a considerable increase in covalent radius from carbon to silicon, but a small increase from silicon to lead. This is due to poor shielding effect of completely filled d and f orbitals.

(ii) Ionisation enthalpy : First ionisation enthalpy is higher than the corresponding members of group 13. Ionisation enthalpy generally decreases down the group. Discrepancies in the trend are also due to poor shielding effect of d and f orbital electrons.

(iii) Electronegativity : Electronegativity from silicon to lead is almost same, but higher than the corresponding group 13 members.

Ø Physical Properties : Boiling point decreases down the group from silicon to lead. Melting and Boiling points are higher than the corresponding group 13 members.Ø Oxidation state : The common oxidation states exhibited by these elements are +4 and +2. Tendency to show +2

oxidation state increases from germanium to lead. Stability of +4 oxidation state decreases down the group. Lead is stable in +2 oxidation state and acts as an oxidising agent in +4 oxidation state.Ø Presence of d-orbital in elements apart from carbon increases their tendency to form complexes by accepting

electron pairs from donors.Ø Chemical reactivity:

(i) Reactivity towards oxygen: Two types of oxides are formed, on heating with oxygen- monoxides and dioxides. Dioxides stend to be more acidic than monoxides. Acidic nature of dioxides decreases down the group; tin and lead dioxide are amphoteric.

(ii) Reactivity towards water: Carbon, silicon and germanium are not affected by water. Tin decomposes steam to form dioxide and dihydrogen gas. Lead is unaffected by water.

76 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI (iii) Reactivity towards halogen : These elements can form halides of formula MX2 and MX4 (where X = F, Cl, Br, I). Stability of dihalides increases down the group, whereas stability of

tetrahalides decreases down the group.Ø Anomalous properties of carbon- Like first member of other groups, carbon also differs from rest of the members

of its group. It is due to its smaller size, higher electronegativity, higher ionisation enthalpy and unavailability of d orbitals.Ø Carbon has a unique property of forming long chains by linking with other carbon atoms. This is called catena-

tion. Tendency of catenation decreases down the group (Lead does not show catenation.)Ø Carbonalsohasuniqueabilitytoformpπ–pπmultiplebondswithitselfandwithotheratomsofsmallsizeand

high electronegativity.Ø Allotropes of carbon :Duetocatenationandformationofpπ–pπmultiplebonds,allotropicformsofcarbon,

namely graphite, diamond and fullerenes exist. (i) Diamond: In diamond, C is sp3 hybridised. (ii) Graphite: In graphite, C is sp2 hybridised. (iii) Fullerenes : In fullerenes, C is sp2 hybridised

Ø Uses of carbon : Carbon, in form of graphite fibres, is used to make tennis rackets and aircrafts. Graphite is used as an electrode in batteries. Activated charcoal is used in water purification and removal of poisonous gases. Dia-mond is used in jewellery and in cutting equipment.Ø SOME IMPORTANT COMPOUNDS OF CARBON AND SILICON—Oxides of Carbon: Two important oxides

of carbon are carbon monoxide, CO and carbon dioxide, CO2.

(i) Carbon Monoxide: Preparation:

(a) 2C(s) + O2 (g) ∆ → 2CO(g)

(b) HCOOH 373

2 4

Kconc H SO. → H2O + CO

(c) C(s) + H2O (g) 473 1273− →Kwater gas

CO(g) + H2(g)

(d) 2C(s) + O2 (g) + 4N2 (g) 1273KProducer gas → 2CO(g) + 4N2(g)

From air Properties: Carbon monoxide is a powerful reducing agent and reduces almost all metal oxides other than those of the

alkali and alkaline earth metals, aluminium and a few transition metals. So, it is used in the extraction of many metals from their oxide ores. e.g.,

Fe2O3(s) + 3CO(g) ∆ → 2Fe(s) + 3CO2(g) ZnO (s) + CO(g) ∆ → Zn(s) + CO2(g) CO molecule (: C ≡ O :) has a lone pair on carbon, therefore acts as a donor and form metal carbonyls. Inhaling CO is very poisonous as it combines with haemoglobin to form complex which is 300 times stable

than oxygen haemoglobin complex, therefore, resulting in loss of oxygen carrying capacity of RBCs. (ii) Carbon Dioxide: Preparation: (a) C(s) + O2 (g) ∆ → CO2(g)

(b) CH4(s) + 2O2 (g) ∆ → CO2(g) + 2H2O (g) (c) CaCO3(s) + 2HCl (aq) → CaCl2(g) + CO2 (g) + H2O(g) (Laboratory preparation) (d) CaCO3(s) → CaO (s) + CO2 (g) (Commercial preparation) Properties: (a) CO2 + H2O → H2CO3 carbonic acid

(b) 6CO2 + H2O h

Chlorophyllv →

C6H12O6 + 6O2 + 6H2O

(Photosynthesis) Unlike CO, it is not poisonous. But increase in CO2 content of the atmosphere may lead to increase

in green house effect and thus, raise the temperature of the atmosphere which might have serious consequences.

Carbon dioxide can be obtained as a solid in the form of dry ice by allowing the liquefied CO2 to expand rapidly. Dry ice is used as a refrigerant for ice-cream and frozen food.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 77 Structure of CO2:

(c) Silicon dioxide, SiO2: Silicon dioxide commonly known as silica, is found in many forms like quartz,

cristobalite and tridymite. Silica is almost non-reactive due to very high Si—O bond enthalpy. However, it is attacked by HF and NaOH.

SiO2 + 2NaOH → Na2SiO3 + H2O SiO2 + 4HF → SiF4 + 2H2O Quartz is used as a piezoelectric material, therefore, used in making extremely accurate clocks, modern

radio and television broadcasting and mobile radio communication. (d) Silicones: Silicones are organosilicon polymers with (-R2SiO3-) repeating units. Silicones have high

thermal stability, high dielectric strength and resistance to oxidation and chemicals. Silicones are used as sealant, greases, electric insulators and for water proofing of fabrics. They are also used in surgical and cosmetic plants as being biocompatible.

(e) Silicates: Silicates exist in nature as feldspar, zeolites, mica and asbestos. Glass and cement are man-made silicates. Silicates have SiO4

4– as basic structural unit in which silicon atom is bonded to four oxygen atoms in tetrahedron fashion.

(f) Zeolites: Replacement of few silicon atom of silicon dioxide by aluminium result in negative charge which is balanced by cations Na+, K+ or Ca2+ to form zeolites. Zeolites are used in cracking and isomerisation of hydrocarbons.

Know the TermsØ Catenation : It is the ability to form compounds in which the atoms are linked to each other in chains or rings.Ø Allotropes : These are elements which exist in two or more different forms in the same physical state.Ø Water Gas : It is a mixture of CO and H2.Ø Product Gas : It is a mixture of CO and N2.

Know the FactsØ Carbon is the 17th most abundant element by mass in the earth’s crust. Ø Carbon compounds are the best example of a group of covalently bonded compounds, studied usually under a

different branch of chemistry called organic chemistry.Ø Diamond: Each carbon is tetrahedrally linked to four other carbon atoms.Ø Graphite: It has a two-dimensional sheet like structure consisting of a number of hexagonal rings fused together.

Graphite conducts electricity along the sheet. It is very soft and slippery.Ø Fullerene: It was discovered collectively by three scientists namely R.E Smalley, R.F Curl and H.W Kroto.

TOPIC-3Group-15 Elements, Properties and Some Important Compounds

Quick ReviewØ GROUP-15 ELEMENTS-The Nitrogen family: It includes Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony

(Sb) and Bismuth (Bi). Nitrogen and phosphorus are non-metals, arsenic and antimony are metalloids while bismuth is a metal.Ø Electronic configuration: They contain five electrons in the outermost shell i.e. a filled s-orbital and three elec-

trons in the p-orbital. (ns2np3). The p-orbital is half-filled, making it more stable.Ø Occurrence: Molecular nitrogen comprises 78% by volume of the atmosphere. In the earth’s crust, it occurs as

sodium nitrate, NaNO3 and potassium nitrate, KNO3.Ø Atomic properties:

(i) Atomic radii—Covalent and ionic radii increase down the group. The lack of regular trend is due to poor shielding effect of d and f orbitals.

(ii) Ionisation enthalpy—Ionisation enthalpy decreases down the group due to increase in atomic size, and is higher than the corresponding group 15 members due to extra stability of half-filled p orbitals.

(iii) Electronegativity—Electronegativity decreases down the group.

78 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XIØ Physical properties:

(i) Metallic character increases down the group. (ii) Boiling point increases upto arsenic and then decreases upto bismuth. (iii) Allotropy is seen in all elements except nitrogen. (iv) Except dinitrogen, all are solid.

Ø Oxidation state: Stability of –3 and +5 oxidation state decreases down the group.Ø Chemical properties:

(i) Reactivity towards hydrogen: All the elements form hydrides of the type EH3 where E = N, P, As, Sb, Bi. The stability decreases from NH3 to BiH3. Reducing character increases down the group. Basic character decreases down the group. Boiling point of NH3 is greater than PH3 because of intermolecular hydrogen bonding. Boiling points increases from PH3 onwards.

(ii) Reactivity towards oxygen: They forms two types of oxides E2O3 and E2O5. The acidic character decreases down the group.

(iii) Reactivity towards halogens: They directly combine with halogens to form trihalides (EX3) and pentahalides (EX5).

(iv) Reactivity towards metals: All the elements react with metals to form their binary compounds exhibiting –3 oxidation state.

Ø Anomalouspropertiesofnitrogen:Nitrogencanformmultiplepπ–pπbonds,whereasphosphorusandarseniccanformdπ–pπbonds.Ndoesnotformpentahalidesbecauseofnon-availabilityofd-orbitals in its valence shell. Ithasabilitytoformpπ–pπmultiplebondswithotherelementshavinghighelectronegativity.Ndiffersfromthe rest of the members of group due to small size, high electronegativity, high ionisation enthalpy and non-availability of d-orbitals.

lDinitrogen (N2) : Preparation: (i) In laboratory: NH4Cl(aq) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl (aq) (ii) By thermal decomposition: Ba(N3)2 → Ba + 3N2 Properties: (i) Dinitrogen is a colourless, odourless, tasteless and non-toxic gas. (ii) N2 has very little reactivity at ordinary temperature. (iii) It forms nitrides with highly electropositive metals.

3Mg + N2 Heat → Mg3N2

6Li + N2 Heat → 2Li3N

lAmmonia (NH3) : Preparation:

(i) In laboratory: 2NH4Cl + Ca(OH)2 Heat → CaCl2 + 2NH3 + 2H2O

(ii) By Haber’s process: N2 (g) + 3H2(g)

Fe/Mo

773K 2NH3(g), H = – 46.1 kJ mol–1

Properties: (i) It is extremely soluble in water. (ii) It is acts as Lewis base Ag+ + 2NH3 [Ag(NH3)2]+

Cu2+ + 4NH3 [Cu(NH3)4]2+

Deep blue Cd2+ + 4NH3 [Cd(NH3)4]2+

(iii) It forms salts with acids ZnSO4(aq) + 2NH4OH(aq) → Zn(OH)2(s) + (NH4)2SO4(aq)

2FeCl3(aq) + 3NH4OH(aq) → Fe2O3.xH2O(s) + 3NH4Cl(aq)

(iv) Reaction with Nessler’s reagent:

2K2[HgI4] + NH3 + 3KOH → [OHg2.NH2]I + 7KI + 2H2O

lNitric Acid (HNO3) :

Preparation: In laboratory: NaNO3 + H2SO4 → NaH2SO4 + HNO3 (conc.) (Brown ppt.)

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 79 (i) By Ostwald’s process: 4NH3 + 5O2 4NO + 6H2O

2NO + O2 → 2NO2

3NO2 + H2O → 2HNO3 + NO

Properties: (i) It is a colourless liquid. (ii) Concentrated nitric acid is a strong oxidising agent. (iii) Reactions : HNO3(aq) + H2O(l) → H3O+(aq) + NO3–(aq)

3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O

(dil)

Cu + 4HNO3 → Cu(NO3)2 + 2NO2 + 2H2O (conc.)

4Zn + 10HNO3 → 4Zn(NO3)2 + 5H2O + N2O (dilute)

Zn + 4HNO3 → Zn(NO3)2 + 5H2O + 2NO2 (conc.)

I2 + 10HNO3 → 2HIO3 + 10NO2 + 4H2O

S8 + 48HNO3 → 8H2SO4 + 48NO2 + 16H2O

P4 + 20HNO3 → 4H3PO4 + 20NO2 + 4H2O

Uses: (i) In the manufacturing of explosives, fertilisers etc. (ii) As a reagent in laboratory. (iii) In nitration of organic compounds. (iv) In preparing aqua-regia.

Ø Oxides of Nitrogen:

S. No.

Formla Name Preparation Properties O.N.

1 N2O Dinitrogen monoxideNH4NO3

Heat → N2O + 2H2OColourless gas, unreac-tive

+1

2 NO Nitrogen monoxide 3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O

Colourless gas, paramag-netic

+2

3 NO2 Nitrogen dioxide2Pb (NO3)2

673K →

2PbO + 4NO2 + O2

Brown gas, reactive, para-magnetic.

+4

4 N2O3 Dinitrogen trioxide 2NO + N2O4 → 2N2O3 Dark blue in liquid or solid state, unstable in the gas phase

+3

5 N2O4 Dinitrogen tetroxide

2NO2

273 K

773 K N2O4

Colourless, exists in equi-librium with NO2 both in the gaseous and liquid state

+4

6 N2O5 Dinitrogen pentoxide 2HNO3 + P2O5 → 2HPO3 + N2O5

Metaphosphoric acid

Unstable as gas; in the solid state

exists as [NO2]+ [NO3]–.

+5

80 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XIØ Structures of Oxides of Nitrogen :

Fig. 1: The molecular and Lewis dot resonance structure of oxides of nitrogen

Ø Allotropes of Phosphorus :

(i) White or yellow phosphorus (ii) Black phosphorus (iii) Red phosphorus

Ø Differences between white and red phosphorus:

Property White Phosphorus Red Phosphorus

State Translucent white waxy solid Iron grey lustrous powder

Odour Garlic Odourless

Physiological action

Poisonous Non-poisonous

Stability Less stable More stable than white P

Solubility Insoluble in water but soluble in CS2. Insoluble in water as well as CS2

Action of air Readily catches fire with greenish glow Does not glow in dark.

Effect of heat Changes to b-black P when heated at 473 K un-der high pressure and changes to red P when heated at 573 K

Changes to a-black P when heated at

803 K in a solid form.

Structure P4 (tetrahedral) Tetrahedral units of P4 joined togetherthrough covalent bond.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 81Ø Phosphine (PH3)

Preparation :

(i) In laboratory: P4 + 3NaOH + 3H2O inert

atmosphereof CO2

∆ PH3 + 3NaH2PO2

(white) (conc.) Sodium hydophosphite Properties :

(i) Colourless gas with rotten fish smell.

(ii) Highly poisonous.

(iii) Pure sample is not spontaneously inflammable.

(iv) Burns in air or oxygen when heated at 150°C.

2PH3 + 4O2 → P2O5 + 3H2O

3CuSO4 + 2PH3 → Cu3P2 + 3H2SO4

3HgCl2 + 2PH3 → Hg3P2 + 6HCl

PH3 + HBr → PH4Br

Ø Phosphorus Trichloride (PCl3)

Preparation :

(i) By passing dry chlorine over heated white phosphorus.

P4 + 6Cl2 → 4PCl3 (ii) By action of thionyl chloride with white phosphorus.

P4 + 8SOCl2 → 4PCl3 + 4SO2 + 2S2Cl2 Properties :

(i) Colourless oily liquid.

(ii) Hydrolyses in the presence of moisture.

PCl3 + 3H2O → H3PO3 + 3HCl

(iii) 3CH3COOH + PCl3 → 3CH3COCl + H3PO3

Shape : Pyramidal in which phosphorus is sp3 hybridised

Ø Phosphorus Pentachloride (PCl5)

Preparation:

(i) Reaction of white phosphorus with excess of dry chlorine.

P4 + 10Cl2 → 4PCl5 (ii) By the action of SO2Cl2 on phosphorus.

P4 + 10SO2Cl2 → 4PCl5 + 10SO2

Properties:

(i) It is yellowish white powder.

(ii) In moist air, it hydrolyses to POCl3 and finally converts to phosphoric acid.

PCl5 + H2O → POCl3 + 2HCl

POCl3 + 3H2O → H3PO4 + 3HCl

(iii) It decomposes on stronger heating

PCl5 Heat → PCl3 + Cl2

(iv) It gives corresponding halides with transition metals

2Ag + PCl5 → 2AgCl + PCl3 Sn + 2PCl5 → SnCl4 + 2PCl3

Shape: Trigonal bipyramidalØ Oxoacids of Phosphorus with structure :

(i) H3PO2 (ii) H4P2O5 (iii) H3PO3 (iv) H4P2O6


Ø The acids which contain P-H bond have reducing characteristics and act as good reducing agent. Example : H3PO2 reduces AgNO3 to Ag. 4AgNO3 + 2H2O + H3PO2 → 4Ag + 4HNO3 + H3PO4

Ø The P-H bonds are not responsible for basicity as they do not ionize to give H+. Only H atoms attached with oxygen are responsible for basicity and are ionisable.

Know the TermsØIndian Salt Petre : It is chemically known as potassium nitrate (KNO3).ØChile Salt Petre : It is chemically known as sodium nitrate (NaNO3).

Know the FactsØ Pnicogens: The nitrogen group is group-15 of the periodic table and is also collectively named the pnicogens or

pnictogens. The word pnicogen is derived from the Greek word pnigein which means ‘to choke or stifle’ which is a property of nitrogen.Ø Fuming nitric acid: Nitric acid containing dissolved NO2 is known as fuming nitric acid. It can be obtained by

distilling concentrated HNO3 with a little of starch.Ø Phosphazenes: These are the cyclic compounds which contain both nitrogen and phosphorus atoms in the al-

ternate position along with two substitutes on each phosphorus atom. These are cyclic trimers, tetramers or polymers in nature.Ø Aqua-regia: Mixture of nitric acid and hydrochloric acid in ratio of 1: 3.

qq


Chapter - 12 : Organic Chemistry-Some Basic Principles and Techniques

TOPIC-1General Introduction, Classification, Nomenclature and Isomerism in Organic Compounds

Quick Review Ø General Introduction: The hydrides of carbon (hydrocarbons) and their derivatives are called organic compounds.

The branch of chemistry which deals with these compounds is called organic chemistry. Ø Berzelius (1808) defined organic chemistry as the chemistry of substances found in living matter and gave the

vital force theory. Urea is the first organic compound synthesised in laboratory, by Wohler. Ø Reasons for Large Number of Organic Compounds (i) Catenation: It is the tendency of self-combination and is maximum in carbon. A carbon atom can combine

with other carbon atoms by single, double or triple bonds. Thus, it forms more compounds than the others. (ii) Tetravalency: Carbon being tetravalent is capable of bonding with four other C atoms or some other

monovalent atoms. Bond angles in sp3 hybridised C is 109°28′, in sp2 hybridised C is 120° and in sp hybridised C is 180°. The C–C, C = C and C ≡ C bond lengths are 1.54 Å, 1.34 Å and 1.20 Å respectively (the shorter the bond,

greater is its strength). Also, greater the bond strength, greater is the bond energy. (iii) Small size: Furthermore, these compounds are exceptionally stable because of the small size of carbon. Ø The three dimensional structure of organic molecule can be represented on paper by using wedge-and-dash

representation. In these formulas, the solid wedge ( ) is used to show a bond coming out of the paper towards the observer and dashed wedge ( ) is used to show a bond going away from the observer.

Ø Classification of organic compound: Organic compounds are broadly classified as follows:

Ø Functional Group: The atom e.g., –Cl, –Br etc., or group of atoms e.g., –COOH, – CHO, which is responsible for the chemical properties of the molecule, is called functional group. Double and triple bonds are also functional groups.

R – F ← functional group Ø Homologous series: Homologous series is defined as a family or group of structurally similar organic compounds

all members of which contain the same functional group, show a gradation in physical and similarity in chemical properties and any two adjacent members of which differ by –CH2 group. The individual members of this group are called homologues and the phenomenon is called homology.

Ø Nomenclature of Organic Compounds: Ø Common name (Common system): Before the IUPAC system of nomenclature, organic compounds were named

after the sources of origin, for example, urea was so named because it was obtained from the urine of mammals. IUPAC (International Union of Pure and Applied Chemistry) System: According to IUPAC system, the name of an organic compound contains three parts: (i) word root, (ii) suffix,

(iii) prefix. (i) Word root: Word root represents the number of carbon atoms present in the principal chain, which is the

longest possible chain of carbon atoms.

84 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI (ii) Suffix: Suffix is of two types, primary suffix and secondary suffix. (a) Primary Suffix: It indicates the type of bond in the carbon chain. (b) Secondary Suffix: Secondary suffix is used to represent the functional group. (iii) Prefix: Prefix is a part of IUPAC name which appears before the word root. Prefix is of two types: (a) Primary prefix: For example, primary prefix cyclo is used to differentiate cyclic compounds. (b) Secondary prefix: Some functional groups are considered as substituents and denoted by secondary

prefixes. For example: Substituted Group Secondary prefix — F Fluoro — Cl Chloro — Br Bromo — NO2 Nitro — NO Nitroso — CH3 Methyl — OCH3 Methoxy Naming of Compounds Containing Functional Groups: The longest chain of carbon atoms containing the

functional group is numbered in such a manner that the functional group is attached at the carbon atoms possessing lowest possible number in the chain. In case of polyfunctional compounds, one of the functional group is selected as principal functional group and the compound is named on that basis. The choice of principal functional group is made on the basis of order of preference.

The order of decreasing priority for the functional group is –COOH, –SO3H, –COOR (R=alkyl group), COCl, –CONH2, –CN, –HC=O, >C=O, –OH, -NH2, >C=C<, –C≡C–. Ø Isomerism: Two or more compounds having the same molecular formula but different physical and chemical

properties are called isomers and this phenomenon is called isomerism. It is of two types- (1) Structural Isomerism: Structural isomerism is shown by compounds having the same molecular formula but

different structural formulae differing in the arrangement of atoms. Structural Isomerism is of four types: (i) Chain isomerism (ii) Position Isomerism (iii) Functional Isomerism (iv) Metamerism (2) Stereoisomerism: When isomerism is caused by the different arrangements of atoms or groups in space,

the phenomenon is called stereoisomerism. The steroeoisomers have same structural formula but differ in arrangement of atoms in space. Stereoisomerism is of two types:

(i) Geometrical or Cis-trans Isomerism (ii) Optical Isomerism

Know the Terms Ø Organic Chemistry: The branch of chemistry which deals with the study of carbon compounds is called organic

chemistry. Ø Isomers: Two or more compounds having the same molecular formula but different properties are called isomers. Ø Aliphatic Compounds: The compounds which have straight or branched chain are called aliphatic compounds. Ø Alicyclic Compounds: The compounds which have carbon atoms joined in the form of a ring are called alicyclic

compounds. Ø Fucntional Group: It may be defined as an atom or group of atoms joined in a specific manner which is responsible

for the characteristic chemical properties of the organic compounds.

Know the Facts Ø Organic compounds may be gases, liquids or solids. Being covalent in nature, these have low boiling point and

melting point and soluble inorganic solvents. These are generally volatile and inflammable. They do not conduct electricity because of the absence of free ions. They possess distinct colour and odour.

Ø Hybridisation affects the bond length and bond enthalpy (strength) in organic compounds. Ø Buckminsterfullerene is a common name given to the newly discovered C60 cluster (a form of carbon) noting its

structural similarity to the geodesic domes popularised by the famous architect R. Buckminster Fuller. Ø Acetic acid is the first organic compound synthesised in the laboratory from its elements.


TOPIC-2Fundamental Concepts in Organic Reaction Mechanism

Quick ReviewØOrganic reactions involve breaking and making of covalent bonds. A covalent bond can undergo fission in two

ways: (i) By Homolytic Fission or Homolysis (ii) By Heterolytic Fission or HeterolysisØHomolytic Fission: In this process, each of the atoms acquires one of the bonding electrons. A-B or The products are called free radicals. They are electrically neutral and have one unpaired electron associated

with them. Homolytic fission is the most common mode of fission in vapour phase. Alkyl radicals are classified as primary, secondary, or tertiary. Alkyl radical stability increases as we proceed from primary to tertiary:

CH < CH CH < CH(CH ) < C(CH )3 2 3 3 2 3 3

Methylfree

radical

Ethylfree

radical

Isopropylfree

radical

Tert-butylfree

radical

ØHeterolytic Fission: In this process, one of atoms acquires both of the bonding electrons when the bond is broken. If one is more electronegative than other which thereby acquires both the bonding electrons and become negatively charged. The products of heterolytic fission are carbonium ions and carbanions.

Ø Carbonium Ions (carbocations): Organic ions which contain a positively charged carbon atom are called carbonium ions or carbocations.

where Z is more electronegative than carbon. The carbocation is sp2-hybridised having trigonal planar geometry. Tertiary carbonium ion is more stable than a secondary, which in turn is more stable than a primary because of inductive effect (+I) associated with alkyl group.

ØCarbanion: Organic ion which contains a negatively charged carbon atom is called carbanions.

where Z is less electronegative than carbon. The carbanions are sp3-hybridised species carrying negative charge and have pyramidal shape. Primary carbanion is more stable than a secondary, which in turn is more stable than a tertiary because of inductive effect (+I) associated with alkyl group.

ØElectrophile: It is a positively charged or neutral species which is electron deficient, i.e. electron seeking e.g., H+, CH3

+, NH4+, AlCl3, SO3, BF3 etc.

86 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Nucleophile: It is negatively charged or neutral species with lone pair of electrons i.e. nucleus seeking e.g.,

OH–, CN–, H2O, R3N, R2NH etc.ØElectron Displacement Effects in Covalent Bonds: Electronic displacements in covalent bonds occur due to the

presence of an atom or group of different electronegativity or under the influence of some outside attacking group. These lead to a number of effects which are as follows:

(i) Inductive effect (ii) Electromeric effect (iii) Resonance or Mesomeric effect (iv) Hyperconjugation. Ø Inductive effect: Electron displacement along the carbon chain in an organic molecule due to the presence of a

polar covalent bond at one end of the chain. Due to different electronegativity, electrons are displaced towards the more electronegative atom. The more electronegative atom acquires a small negative charge (δ). The less electronegative atom acquires a small positive charge (δ).

Ø Atoms or groups which lose electrons towards a carbon atom are said to have a + I effect. Those atoms or roups which draw electrons away from a carbon atom are said to have a -I effect. Some common atoms or groups which cause +I or -I effects are shown below:

(i) The – I effect (electron attracting) shows the trend: NO2 > F > COOH > Cl > Br > I > OH > C6H5. (ii) The + I effect (electron releasing) shows the trend: (CH3)3 C— > (CH3)2CH— > CH3CH2 — > CH3. Ø Electromeric Effect (E effect): It is a temporary effect. It is defined as the complete transfer of a shared pair of

p-electrons to one of the atoms joined by a multiple bond on the demand of an attacking reagent. The shifting of the electrons is shown by a curved arrow ( ) . There are two distinct types of electromeric effect:

(i) Positive Electromeric Effect (+E effect):Inthiseffect,theπ-electronsofthemultiplebondaretransferredtothat atom to which the reagent gets attached. For example,

(ii) Negative Electromeric Effect (–E effect): In this effect, the p - electrons of the multiple bond are transferred to that atom to which the attacking reagent does not get attached. For example,

Ø Resonance or Mesomeric Effect: Thepolarityproducedinthemoleculebythe interactionof twoπ-bondsorbetweenaπ-bondandalonepairofelectronspresentonanadjacentatom.Therearetwotypesofresonanceormesomeric effects designated as R or M effect.

Positive Resonance Effect (+R effect): Those atoms which lose electrons towards a carbon atom are said to have a +M effect or +R effect. For example: —Cl, —Br, —I, —NH2, —NR2 , —OH, —OCH3.

Negative Resonance Effect (-R effect): Those atoms or groups which draw electrons away from a carbon atom are said to have a -M effect or -R effect. For example: –COOH , -CHO , -CN

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 87 Ø Hyperconjugation: Orbital interactions between the p systems and the adjacent sigma bonds of the substituent

groups (s) in organic compounds. It involves delocalisation of s electrons of C—H bond of an alkyl group directly attached to an atom of unsaturated

system or to an atom with an unshared p orbital. The s electrons of C—H bond of the alkyl group enter into partial conjugation with the attached unsaturated system or with the unshared p orbital. Hyperconjugation is a permanent effect which is also called “No bond resonance”.

In general, greater the number of alkyl groups attached to a positively charged carbon atom, the greater is the hyperconjugation.

Ø Types of Organic Reactions and Mechanisms: (i) Substitution reactions—Reactions involving replacement of an atom or group in a molecule by different

atoms or groups e.g., chlorination of methane. (ii) Addition reactions—Reactions when two molecules form a single product e.g., reaction of but-1-ene with

HBr to give 2-bromobutane. (iii) Elimination reactions—Reactions in which two atoms or groups in a molecule are eliminated to give the

product e.g., formation of alkene from alkyl bromide. (iv) Rearrangement reaction—Reactions involving migration of atoms to other position within the molecule.

Know the TermsØCarbocation: A species having a carbon atom possessing sextet of electrons and a positive charge is called

carbocation.ØCarbanion: A species having a carbon atom possessing sextet of electrons and an unshared pair of electrons and

negative charge is called carbanion. ØIonic reactions: The organic reactions which proceed through heterolytic bond cleavage are called ionic or

heteropolar reactions. ØFree radical reactions: Organic reactions which proceed by homolytic fission are called free radical or homopolar

or non-polar reactions.ØResonance energy: The difference in energy between the actual structure and the lowest energy resonance

structure is called resonance energy.

Know the FactsØHeterolytic and homolytic bond fission results in the formation of short-lived fragments called reaction

intermediates. ØDuring a polar organic reaction, nucleophile attacks an electrophilic centre of the substrate which is that specific

atom or part of the elecrophile that is electron deficient.ØThe energy of actual structure of the molecule (the resonance hybrid) is lower than that of any of the canonical

structures.ØThe resonance structures have (i) the same positions of nuclei and (ii) the same number of unpaired electrons.


TOPIC-3Methods of Purification of Organic Compound

Quick Review Ø Purification of organic compounds can be done by sublimation, crystallisation, distillation, differential extraction

and chromatography. Sublimation: This technique is used only for those solids which changes from solid to vapour state directly. Crystallisation: This is used for purification of solid organic compounds. This is based on the difference in the

solubilities of compound and the impurities in a suitable solvent. Distillation: It can be used (i) to separate volatile liquids from non-volatile impurities (ii) to separate the liquids

having sufficient difference in their boiling points (>25K), liquids having different boiling points vaporise at different temperatures. The vapours are cooled and the liquids so formed are collected separately.

Fractional distillation: It is used for separating liquids if the difference in boiling points is not much. In this technique, vapours of a liquid mixture are passed through a fractionating column before condensation. A fractionating column provides many surfaces for heat exchange between the ascending vapours and the descending condensed liquid. This technique is used to separate different fractions of crude oil in petroleum industry.

Distillation under reduced pressure: It is carried out to purify those liquids that have high boiling points and those which decompose at or below their boiling points.

Steam distillation: It is used to purify those substances which are steam volatile and immiscible with water. In steam distillation, the liquids boils when the sum of vapour pressures due to the organic compound (P1) and that due to water (P2) becomes equal to the atmospheric pressure (P), i.e., P = P1 + P2.

Differential extraction: Here extraction of compound takes place based on difference in solubility. An organic compound present in aqueous medium is separated by shaking it with an organic solvent in which it is more soluble than in water. Organic solvent and water being immiscible are separated by using separating funnel. Removal of organic compound by distillation or evaporation yields pure compound.

Chromatography: It is based on the general principle of distributing the components of a mixture of organic compounds between two phases-a stationary phase and mobile phase. The stationary phase can be solid or liquid supported on a solid, while the mobile phase is a liquid or a gas. When the stationary phase is a solid, the basis of separation is adsorption; when it is a liquid, the basis of separation is partition.

Ø Qualitative Analysis of Organic Compounds: Almost all the organic compounds contain carbon and hydrogen (CCl4 is exception as it does not contain hydrogen). They may contain oxygen, nitrogen, sulphur, halogens and phosphorus.

Detection of Carbon and Hydrogen: Carbon and hydrogen are detected as carbon dioxide and water respectively by heating with cupric oxide. Carbon dioxide is detected as it turns lime water milky, however, water is detected as it turns anhydrous copper sulphate to blue.

C + 2CuO ∆ → 2Cu +CO2

2H + CuO ∆ → Cu + H2O CO2 + Ca(OH)2 → CaCO3 ↓+ H2O 5H2O + CuSO4 → CuSO4.5H2O White Blue

Lassaigne‘s test: Nitrogen, sulphur, halogens and phosphorus present in organic compound are detected by ‘Lassaigne‘s test. In this test, organic compound is fused with sodium metal so that elements present in organic compound are converted from covalent form into the ionic form.

Na + C + N ∆ → NaCN 2Na + S ∆ → Na2S Na + X ∆ → NaX (X = Cl, Br or I)

These fused mass are extracted by boiling with water to form sodium fusion extract. Test for Nitrogen—The sodium fusion extract is boiled with iron (II) sulphate and then acidified with

concentrated sulphuric acid. The formation of Prussian blue colour confirms the presence of nitrogen.

6CN– + Fe2+ → [Fe (CN)6]4–

3[Fe (CN)6]4– + 4Fe3+ xH O2 → Fe4[Fe(CN)6]3.xH2O Prussian blue (Ferric ferrocyanide)

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 89 Test for Sulphur—Add 2-3 drops of freshly prepared sodium nitroprusside solution to sodium metal extract.

Appearance of violet colour ceonfirms the presence of sulphur.

[Fe (CN)5NO]2– + S2– → [Fe(CN)5NOS]4–

Sulphonitropruside ion (violet colour) If both sulphur and nitrogen are present, sodium thiocyanate is produced in sodium metal extract which gives

red colour with Fe3+ ions while testing for nitrogen.

3NaSCN + Fe3+ → Fe(SCN)3 + 3Na+

Ferric thiocyanate (Blood red)

Test for Halogens—For this, sodium fusion extract is acidified with nitric acid and then treated with silver nitrate: (i) A white precipitate, soluble in ammonium hydroxide shows the presence of chlorine. (ii) A yellowish precipitate, sparingly soluble in ammonium hydroxide shows the presence of bromine. (iii) A yellow precipitate, insoluble in ammonium hydroxide shows the presence of iodine.

X– + Ag+ → AgX (X = – Cl, – Br, – I)

Test for Phosphorus—Compound is heated with an oxidising agent (e.g., sodium peroxide), that oxidises phosphorus to phosphate. This solution on boiling with nitric acid followed by treatment with ammonium molybdate produces canary yellow precipitate confirming phosphorus:

Na3PO4 + 3HNO3 → H3PO4 + 3NaNO3 H3PO4 + 12(NH4)2MoO4 + 21HNO3 → (NH4)3PO4.12MoO3 + 21NH4NO3 + 12 H2O Ammonium molybdate Ammonium phosphomolybdate (Canary yellow precipitate) Ø Quantitative analysis: Quantitative analysis of organic compound gives the percentage composition of elements

in an organic compound. Ø Estimation of Carbon and Hydrogen: Let the mass of organic compound be m g, mass of water and carbon

dioxide produced be m1 and m2 g respectively

Percentage of carbon= 12×m ×100

44×m2

Percentage of hydrogen= 2×m ×100

18×m1

Ø Estimation of Nitrogen: (i) Dumas method: Let the mass of organic compound = m g Volume of nitrogen collected = V1 mL Room temperature = T1K

Volume of nitrogen at STP=

p ×V ×273760×T

1 1

1

Percentage of nitrogen = 28×volomeof N atSTP×100

22400×weight of organiccompound2

(ii) Kjeldahl’s method: Volume of acid of normality N1 neutralised by NH3= V1cm3

1000 cm3 of 1N NH3=17 g of NH3 V1 cm3 of N1 acid= 14 g of nitrogen

= 14×N ×V

10001 1 g of Nitrogen

Percentage of nitrogen= 14×N ×V

1000×

100W

1 1

Where N1 is the normality of acid, V1= volume of acid and W= weight of organic compound. Estimation of Halogens: Let the mass of organic compound taken = m g Mass of AgX formed = m1 g 1 mol of AgX contains 1 mol of X Mass of halogen in m1g of AgX

= Atomicmassof X×m gmolecular massof AgX

1

Percentage of halogen = Atomicmassof X×m g×100molecular massof AgX×m

1

Estimation of Sulphur: Let the mass of organic compound taken = m g

90 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI and the mass of barium sulphate formed = m1 g 1 mol of BaSO4 = 233 g BaSO4 = 32 g sulphur

m1 g BaSO4 contains 32×m

2331 g of sulphur

Percentage of sulphur = 32×m ×100

233×m1

Estimation of Phosphorus: Let the mass of organic compound taken = m g and mass of ammonium phosphomolydate = m1g Molar mass of (NH4)3PO4.12MoO3 = 1877 g

Percentage of phosphorus= 31×m ×100

1877×m1 %

If phosphorus is estimated as Mg2P2O7,

Percentage of phosphorus= 62×m ×100222×m

1 %

Know the Terms Ø Chromatography: It is a technique used to separate mixtures into their components, purify compounds and also

to test the purity of compounds. Ø Adsorption Chromatography: It is a chromatographic technique which is based on the fact that different

compounds are adsorbed on an adsorbent at different degrees. Ø Column Chromatography: It is a technique which involves separation of a mixture over a column of adsorbent

(stationary phase) packed in a glass tube. Ø Thin Layer Chromatography: It is a technique which involves separation of substances of a mixture over a thin

layer of an adsorbent coated on glass plate. Ø Rf Value: The relative adsorption of each component of the mixture is expressed in terms of its retardation factor

i.e. Rf value. Ø Partition Chromatography: It is a technique which is based on continuous differential partitioning of components

of a mixture between stationary and mobile phases.

Know the Facts Ø Mixed melting point: The melting point of two thoroughly mixed substances is called mixed melting point. This

can also be used for ascertaining the purity of a compound. Ø A number of chemical methods have also been used to separate a mixture of organic compounds. These methods

are based upon the distinguishing chemical properties of one class of organic compounds from the others. Ø Criteria of purity of organic compounds: The purity of an organic compound can be ascertained by determining

its some physical constants like m.p., b.p., specific gravity, refractive index and viscosity.

qq

Chapter - 13 : Hydrocarbons

TOPIC-1Saturated Hydrocarbons (Alkanes)

Quick Review Ø Classification of Hydrocarbons: Hydrocarbons are the compounds of carbon and hydrogen only. Hydrocarbons

are mainly obtained from coal and petroleum, which are the major sources of energy. Hydrocarbons are classified as saturated hydrocarbon, unsaturated hydrocarbon and aromatic hydrocarbon.

A hydrocarbon is said to be saturated if it contains only C—C single bonds. For example: Ethane CH3 —CH3. Unsaturated hydrocarbons contain carbon-carbon multiple bonds i.e., double bonds, triple bonds or both. For

example: CH2=CH2.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 91 Aromatic hydrocarbon compounds consists of at least one aromatic ring. For example:

Ø Alkanes : Alkanes are the simplest organic compounds made of carbon and hydrogen only. They have the general formula CnH2n+2 (where n = 1, 2, 3 etc.) Alkanes contain strong C —C and C —H bonds. Therefore, this class of hydrocarbons is relatively chemically inert. Hence, they are sometimes referred to as paraffins (Latin : parum affinis = little affinity).

Ø Methane (CH4) is the first member of this family. In methane, carbon forms single bonds with four hydrogen atoms. All H—C—H bond angles are of 109.5°. Methane has a tetrahedral structure. C—C and C—H bonds are formed by head-on overlapping of sp3 hybrid orbitals of carbon and 1s orbitals of hydrogen atoms.

Ø Alkanes exhibit chain isomerism, position isomerism and conformational isomerism. Ø Preparation of Alkanes : (i) Alkanes can be prepared by hydrogenation of unsaturated hydrocarbons in the presence of catalyst (Ni, Pd or

Pt).

CH2 = CH2 + H2 Ni Pd Pt/ / → CH3—CH3 Ethene Ethane

HC ≡ CH + 2H2 Ni Pd Pt/ / → H3C—CH3

(ii) By reduction of alkyl halides- Alkanes can be prepared by the reduction of alkyl halides (except fluorides) with zinc and dilute hydrochloric acid.

CH3—Cl + H2 Zn H, + → CH4 + HCl

Chloromethane Methane

(iii) Wurtz Reaction— This reaction is used to increase the length of the carbon chain.

CH3Br + 2Na + BrCH3 dry ether → CH3—CH3 + 2NaBr

(iv) By sodalime- This reaction is used for descending of series as the alkane obtained has one carbon less than the parent compound. CaO is more hygroscopic than NaOH and it keeps NaOH in dry state.

CH3COO–Na+ + NaOH CaO∆

→ CH4 + Na2CO3

(v) Kolbe’s Electrolytic Method— 2CH3COO– Na+ + H2O → CH3—CH3 + 2CO2 + H2 + 2NaOH Sodium acetate Ethane Ø Physical properties : With the increase of branching, the boiling points of alkanes decrease as on increasing the

branching, the molecule attains the shapes of a sphere, decreasing surface area, thereby decreasing van der Waals forces. Alkanes with even number of carbon atoms have higher melting points as compared to next higher or lower alkanes with odd number of carbon atoms.

Ø Chemical Properties : (i) Substitution Reactions: One or more atom of alkanes is replaced by halogens, nitro group and sulphonic

acid group, e.g., halogenation.

CH4 + Cl2 hv → CH3Cl + HCl

(ii) Combustion—Due to the evolution of a large amount of heat during combustion, alkanes are used as fuels.

CnH2n + 2 +3

2n+1

O2 → nCO2 + (n + 1) H2O

e.g., C4H10 (g) + 13/2 O2(g) → 4CO2(g) + 5H2O (l);∆H°=–2875·84kJmol–1

(iii) Controlled Oxidation : (i) 2CH4 + O2 Cu K atm/ /523 100 → 2CH3OH Methanol

(ii) CH4 + O2 Mo O2 3

∆ → HCHO + H2O

Methanal

2CH3CH3 + 3O2 CH COO Mn3 2( ) → 2CH3COOH + 2H2O

Ethanoic acid


(iv) (CH3)3CH KMnO

O4

[ ] → (CH3)3COH 2-Methylpropane 2-Methylpropan-2-ol (3°-alkanes) (iv) Isomerisation—

CH3(CH2)4CH3 Anhyd AlCl HCl. / 3 →

CH3

CH CH (CH ) CH2 2 2 3

2-Methylpentane

+CH CH CH CH CH3 2 2 3

CH3

3-Methylpentane

(v) Aromatization or Reforming—Dehydrogenation and cyclisation to benzene and its homologues.

(vi) Reaction with steam—(Industrial preparation of dihydrogen gas)

CH4 + H2O Ni∆

→ CO + 3H2

(vii) Pyrolysis—Decomposition of higher alkanes to lower alkanes by the application of heat is called pyrolysis or cracking.

C H6 14

773 K

C H + H6 12 2

C H + C H4 8 2 6

C H + C H + CH3 6 2 4 4

Ø Conformations of Alkanes : Alkanes have C-C sigma (s) bonds and rotation about C-C single bond is allowed. This rotation results in different spatial arrangements of atoms in space which can change into one another, such spatial arrangements are called conformations or conformers or rotamers. Alkanes can have infinite number of conformations. Ethane shows eclipsed, staggered and skew conformations. Eclipsed form is least stable but staggered form is most stable due to greater distance between the bond pairs or lesser torsional strain. The energy difference between the two extreme forms is of the order of 12.5 kJ mol–1.

Know the Terms Ø Cycloalkanes : Cyclic hydrocarbons having molecular formula CnH2n. These are isomeric with alkenes. Ø Cracking : The process of breaking down less volatile higher molecular mass hydrocarbons from petroleum into

different types of more volatile lower molecular mass hydrocarbons by heating in the presence of catalyst. Ø Reforming : Cyclization followed by aromatization (or dehydrogenation of 6-8 C alkanes on heating to 873 K in

the presence of Pt catalyst (e.g., n-heptane → methyl cyclohexane → toluene). Ø Knocking : A sharp metallic rattling sound produced in an internal combustion engine. Ø Octane Number : A scale for determining knocking quality in petrol. n-Heptane with very high knocking has

octane number 0 while iso-octane has maximum value 100. Ø Cetane Number : A scale for determining knocking quality of a diesel fuel. Cetane number of n-hexadecane

(cetane) is 100 and that of a-methylnapthalene is zero.

Know the Facts Ø Iodination is a reversible reaction. So it is carried out by heating alkane in the presence of oxidising agent like iodic

acid (HIO3) or nitric acid (HNO3) or mercuric oxide (HgO) which oxidises HI formed during the reaction. Ø Fluorination of alkane takes place explosively resulting even in the rupture of C—C bond in higher alkanes. Ø Petroleum is an important source of aliphatic hydrocarbons. It is refined in to several useful fractions and is a

source of several petrochemicals. Coal tar is a rich source of aromatic hydrocarbons. Ø LPG and CNG are petroleum products used as energy source in automobile industry and domestic fuel. Natural

gas contains about 90% CH4.


TOPIC-2Unsaturated Hydrocarbons (Alkenes and Alkynes)

Quick Review Ø Alkenes : These are unsaturated non-cyclic hydrocarbons containing at least one double bond and which have

sp2 -hybridisation with 120° bond angle. Alkenes are also called olefins [oil forming] which indicates their high reactive nature. Alkenes have general formula CnH2n, where n = 2,3,4 …Example- C2 H4 (ethene), C3H6 (propene), etc.

Ø In ethylene, the carbon atoms are attached to each other by a weak p bond and a strong σ bond. The bond results from the overlap of two sp2 hybrid orbitals. The p bond is formed from overlap of the unhybridised p-orbitals. Ethylene is a planar molecule.

Ø Isomerism : Alkenes show chain, position and geometrical isomerism. Chain isomers have different chains of carbon atoms, e.g.,

CH2 = CH—CH2—CH3 and CH CH CH2 3

CH3

Position isomers differ in the position of double bonds.

CH2 = CH—CH2—CH3 and CH3—CH = CH—CH3

The stereoisomerism is exhibited by alkenes due to difference in the spatial arrangement of groups around double bonded carbon atoms. The simplest alkene that can exhibit geometrical isomerism is but-2-ene. Cis form of alkene is more polar than the trans form.

cis-But-2-ene trans-But-2-ene

H

C C

H C3 CH3

H H

C C

H C3

CH3

H

The boiling point of cis-form is more than the trans-form due to high polarity of cis-form. However, m.p. of trans-form is more than that the cis-form.

Ø Preparation of Alkenes : (i) From alkynes:

RC ≡ CR′+ H2 Pd C S or quinolineLindlar s catalyst/

'−

( ) →

C C

R

H

R'

Hcis-Alkene

RC CR' + H2

Na/liquid NH3 C C

R

HR'

H

trans-Alkene

(ii) From alkyl halides : (Dehydrohalogenation) The rate of the reaction : for halogens is : iodine > bromine > chlorine for alkyl group is : tertiary > secondary > primary

H

C

H

H C3

(X = Cl, Br, I)

H

C

X

��

Halc. KOH

�

C C

H

H

H

H

(iii) From vicinal dihalides : (Dehalogenation) CH3CHBr—CH2Br + Zn → CH3CH = CH2 + ZnBr2

94 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI (iv) From Alcohols : (Acidic dehydration)

Ø Physical Properties : Alkene as a class resembles alkanes in physical properties, except in types of isomerism and

difference in polar nature. Alkenes show a regular increase in boiling point with increase in size. Ø Chemical properties :Alkenesundergoelectrophilicadditionreactionbecauseoflooselyheldπ-electrons. (i) Addition of dihydrogen :

H2C = CH2 H

Ni Pt Pd2

/ / → H3C—CH3

Ethene Ethane (ii) Addition of Halogen : This reaction is used for the test of unsaturation.

H C = CH + Br Br2 2

CCl4

CH CH2 2

Br Br

(Reddish orange) (Colourless)

(iii) Addition of hydrogen halides : CH2 = CH2 + HBr → CH3—CH2—Br (symmetrical alkene) However, addition of hydrogen halide to an unsymmetrical alkene takes place according to Markovnikov’s

rule. This rule states that, “the negative part of the addendum (adding molecule) gets attached to that carbon atom which possesses lesser number of hydrogen atoms”. e.g.,

Peroxide effect or Kharash effect or Anti-Markovnikov addition : In the presence of peroxide, addition of HBr to unsymmetrical alkene takes place contrary to Markovnikov rule. This reaction is known as peroxide or Kharash effect or anti-Markovnikov addition. e.g.,

H C CH CH + HBr3 2

1-Bromopropane

CH CH CH Br3 2 2

(C H CO) O6 5 2 2

This reaction proceeds by free radical mechanism. Peroxide effect is observed only in case of HBr and not in case of H—Cl and H—I.

(iv) Addition of sulphuric acid :

O

S

O

CH = CH + H2 2

Ethyl hydrogen sulphate

O HO CH —CH —OSO —OH (C H HSO )3 2 2 2 5 4

(v) Addition of water : In presence of few drops of concentrated sulphuric acid.

(vi) Oxidation : (a) Oxidation with cold dilute, aqueous solution of potassium permanganate (Baeyer’s reagent). This

reaction is used as a test for unsaturation as it decolourises KMnO4 solution.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 95 (b) Oxidation with acidic potassium permanganate or acidic potassium dichromate to ketones and/or acids. (CH3)2C = CH2 KMnO H4 / +

→ (CH3)2C = O + CO2 + H2O 2-Methylpropene Propan-2-one

CH3—CH = CH—CH3 KMnO H4 / + → 2CH3COOH

But-2-ene Ethanoic acid (vii) Ozonolysis :

(viii) Polymerisation :

n(CH CH = CH )2 2 —(—CH —CH —) —2 2 n

Polythene

High Temp. (Pressure)

Catalyst

Ø Alkynes : Alkynes are unsaturated hydrocarbons that contain at least one carbon-carbon triple bond. General formula of alkyne is CnH2n-2 e.g., C2H2 (ethyne), C2H4 (propyne). The first and the most important member of this series of hydrocarbons is acetylene, C2H2 and hence they are also called the Acetylenes.

In ethyne, the carbon atoms are sp hybridised. They are attached to each other by a s-bond and two p-bonds. The s -bond results from the overlap of two sp hybrid orbitals. The p bonds are formed from the separate overlap of the two p-orbitals from the two adjacent carbon atoms. The other sp hybrid orbital of each carbon atom forms a p bond with another carbon or hydrogen atom. Ethyne is a linear molecule.

Ø Preparation of alkynes : (i) From calcium carbide : Ethyne is prepared by treating calcium carbide with water. Calcium carbide is

prepared as follows:

CaCO3 ∆ → CaO + CO2

Limestone Quicklime

CaO + 3C → CaC2 + CO Calcium carbide CaC2 + 2H2O → Ca(OH)2 + C2H2 Acetylene (ii) From vicinal dihalides : When vicinal dihalides react with alcoholic potassium hydroxide then they

undergo dehydrohalogenation. One molecule of hydrogen halide is eliminated to form alkenyl halide which on treatment with sodamide gives alkyne.

Ø Physical Properties : Physical properties of alkynes follow the same trend of alkenes and alkanes. Alkynes are insoluble in water but soluble in organic solvents like ethers, carbon tetrachloride and benzene. Melting point, boiling point and density increase with increase in molar mass.

Ø Chemical Properties : Alkynes show electrophilic as well as nucleophilic addition reactions. (i) Acidic character of alkyne : HC ≡ CH + Na → HC ≡ C–Na+ + ½H2 Monosodium ethynide HC ≡ C–Na+ + Na → Na+C– ≡ C–Na+ + ½H2 Disodium ethynide CH3—C ≡ C—H + Na+NH2

– → CH3C ≡ C–Na+ + NH3 Sodium propynide

This reaction can be used for the distinction among alkynes, alkenes and alkanes and even for distinguishing terminal alkynes from non-terminal alkynes as only terminal alkynes undergo this reaction.

96 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI (ii) Addition reactions :

Alkynes undergo electrophilic addition reaction and addition to unsymmetrical alkynes takes place according

to Markovnikov’s rule. The addition product formed depends upon the stability of vinylic cation. (a) Addition of dihydrogen :

CH3—C ≡ CH + H2 Ni Pd Pt/ / → [CH3—CH = CH2] H2 → CH3—CH2—CH3

Propyne Propene Propane

(b) Addition of halogens :

This reaction is used as a test of unsaturation as reddish orange colour of CCl4 is decolourised. (c) Addition of hydrogen halides :

(d) Addition of water :

(e) Polymerisation : (1) Linear polymerisation : Acetylene → Polyacetylene (2) Cyclic polymerisation :

Know the TermsØ Alkenes: The unsaturated hydrocarbons containing atleast one double bond are called alkenes. Ø Alkynes: The unsaturated hydrocarbon containing atleast one triple bond are called alkynes. Ø Cis-isomer: The geometrical isomer in which two identical atoms or groups lie on the same side of the double

bond is called cis-isomer. Ø Trans-isomer: The geometrical isomer in which two identical atoms or groups lie on the opposite sides of the

double bond is called trans-isomer. Ø Acidic Dehydration of Alcohols: The reaction in which a water molecule is eliminated from the alcohol molecule

in the presence of an acid is called acidic dehydration of alcohols. Ø Polymerisation: The process by which a long chain like molecule is obtained by the intermolecular combination

of small molecules (monomers) is called polymerisation.

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 97Know the Facts Ø Peroxide effect is applicable only to HBr and not to HF, HCl and HI. Addition of HF, HCl and HI takes place

according to Markovnikov’s rule even in the presence of peroxide. Ø The alkaline potassium permanganate solution is known as Baeyer’s reagent. It has bright pink colour. It oxidises

alkenes to glycols which is colourless. This reaction is used as a test for the presence of double bond in a molecule. This is also known as Baeyer test.

Ø Bromine water test and Baeyer’s test are used to detect the presence of double bond while ozonolysis is used to detect the position of double bond.

TOPIC-3Aromatic Hydrocarbons: Benzene

Quick Review Ø These hydrocarbons are also known as ‘arenes’. Most of the aromatic hydrocarbons were found to contain

benzene ring. Aromatic compounds containing benzene ring are known as benzenoids and those not containing a benzene ring are known as non-benzenoids. Benzene undergoes substitution more readily than addition inspite of the presence of high degree of unsaturation due to highly resonance stabilized benzene ring. Benzene molecule is a planar or flat molecule in which all the carbon atoms are sp2 hybridised. It has hexagonal ring of six carbon atoms with three double bonds at alternate positions. It is resonance stabilised and the structure may be represented as given below.

(A) (B) (C)

or

Ø Aromaticity : It is a property of the sp2 hybridised planar rings in which the p orbitals allow cyclic delocalization ofπelectrons.AromaticcompoundshavespecificelectronicstructureinaccordancewithHuckelrule.Therulestatesthatallplanarcyclicconjugatedpolyenescontaining(4n+2)πelectronswheren=0,1,2,…arearomaticin nature.

Ø Preparation of benzene : (i) Cyclic polymerisation of ethyne:

(ii) Decarboxylation of Aromatic acids:

CaO+ NaOH

�

+ Na CO2 3

COONa

(iii) Reduction of Phenol:

�

+ Zn + ZnO

OH

Ø Physical Properties : Benzene itself is a good solvent for many organic and inorganic substances e.g., fat, resins, sulphur and iodine. It burns with a luminous, sooty flame in contrast to alkanes and alkenes which usually burn with a bluish flame.

Ø Chemical Properties : Benzene undergoes following types of chemical reactions: (i) Electrophilic Substitution Reaction (ii) Addition Reaction Electrophilic Substitution Reactions : Aromatic hydrocarbons undergo electrophilic substitution reactions,

e.g., nitration, halogenation, sulphonation and Friedel-Craft alkylation/acylation.

98 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI (i) Nitration :

(ii) Halogenation :

(iii) Sulphonation :

+ H SO (SO )2 4 3 + H O2

�

Benzene sulphonicacid

SO H3

Fuming sulphuricacid

(iv) Friedel-Crafts alkylation reaction:

(v) Friedel-Craft’s acylation reaction:

+ CH COCl3 + HCl

Acetophenone

COCH3

Anhyd. AlCl3

Acetyl chloride

�

Benzene on treatment with excess of chlorine in the presence of anhydrous AlCl3 can be chlorinated to hexachlorobenzene (C6Cl6)

+ 6Cl2

Anhyd. AlCl3

�

+ 6HCl

Hexachlorobenzene

(C Cl )6 6

Cl

Cl

Cl Cl

Cl Cl

Addition reactions :

+ 3H2

Ni

High temp. & pressure

Cyclohexane

+ 3Cl2

UV

500 K

Benzene hexachloride(BHC or gammaxane)

Cl

Cl

Cl Cl

Cl Cl

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 99 Combustion—Benzene burns with sooty flame : C6H6 +15/2 O2 → 6CO2 + 3H2O Ø The presence of substituent groups in aromatic ring activates/deactivates electrophilic substitution and also

directs the orientation of incoming groups. Electron donating substituents facilitate further substitution with incoming groups at o- and p-positions while electron attracting groups deactivate the ring with substitution at m-position.

– CH3, – OH, – OR, – NH2, – NHR, – NR2 are electron donating (or activating groups) while – NO2, – CN, – CHO, – COOH, – SO3H are electron withdrawing (or deactivating groups). Halogens (– Cl, – Br, – I) are o-, p-directing but moderately deactivating (due to strong – I effect).

Ø Polynuclear aromatic hydrocarbons (PAHs) have a number of condensed benzene rings and are suspected carcinogens (toxic and they are having cancer producing property) They are actually formed due to incomplete combustion of some organic materials like tobacco, coal, petroleum etc

Know the TermsØ Arenes: Aromatic hydrocarbons are known as arenes. Ø Benzenoids: Aromatic compounds containing benzene ring are called benzenoids. Ø Non-benzenoids: Aromatic compounds which do not contain benzene ring are called non-benzenoids. Ø Aromaticity: It is a property of conjugated cycloalkenes in which the stabilization of the molecule is enhanced

due to the ability of the electrons in the p orbitals to delocalize. This acts as a framework to create a planar molecule.

Know the Facts Ø Originally fragrant substances were called aromatic compounds. Ø In the laboratory, benzene was first prepared by heating benzoic acid or phthalic acid with calcium oxide. Ø Most of the benzene is produced from petroleum. Ø Compounds having atoms other than carbon in a ring system, which satisfies the conditions of aromaticity are

called heteroaromatics.

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Chapter - 14 : Some Basic Concepts of Chemistry

TOPIC-1Atmospheric Pollution

Quick Review Ø Environmental Chemistry: The branch of chemistry which deals with the study of various chemical changes and

reactions occurring in environment is called Environmental chemistry. It includes our surroundings as air, water, soil, forest etc.

Ø Environmental Pollution: Environmental pollution is the effect of undesirable changes in our surroundings that have harmful effects on plants, animals and human beings.

Ø Pollutants: A substance which causes pollution is known as pollutant. Pollutants can be solid, liquid or gaseous substances, present in higher concentration. It can be produced due to human activities or natural happenings.

Ø Atmospheric Pollution: The atmosphere that surrounds the earth is not of the same thickness at different heights. Atmospheric pollution is generally studied as tropospheric and stratospheric pollution.

Ø Tropospheric Pollution: Tropospheric pollution occurs due to the presence of undesirable solid or gaseous particles in the air. The following are the major gaseous and particulate pollutants present in the troposphere:

l Gaseous air pollutants: These are oxides of sulphide, hydrocarbons, ozone and other oxidants. l Particulate pollutants: These are dust, mist, fumes, smoke, smog etc. (i) Gaseous air pollutants: Oxides of Sulphur: Oxides of sulphur are produced when sulphur containing fossil fuel is burnt. The most

common species sulphur dioxide is a gas that is poisonous to both animals and plants.

2SO2(g) + O2(g) → 2SO3(g)

SO2(g) + O3(g) → SO3(g) + O2(g)

SO2(g) + H2O2(l) → H2SO4(aq)

100 ] Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI Oxides of Nitrogen: NO2 is oxidised to nitrate ion. When fossil fuel is burnt, dinitrogen and dioxygen combine to

yield significant quantities of nitric oxide(NO) and nitrogen dioxide (NO2).

N2(g) + O2(g) → 2NO(g) [at 1483 K]

2NO(g) + O2(g) → 2NO2(g)

NO(g) + O3(g) → NO2(g) + O2(g)

Hydrocarbons: Hydrocarbons are carcinogenic, i.e., they cause cancer. They harm plants by causing ageing, breakdown of tissues and shedding of leaves, flowers and twigs.

Oxides of Carbon: l Carbon monoxide: Carbon monoxide is one of the most serious air pollutants. It is highly poisonous to living

beings because of its ability to block the delivery of oxygen to the organs and tissues. It binds to haemoglobin to form carboxy haemoglobin, which is about 300 times more stable than the oxygen-haemoglobin complex. This results in headache, weak eyesight, nervousness and cardiovascular disorder.

l Carbon dioxide: Carbon dioxide is released into the atmosphere by respiration, burning of fossil fuels for energy and by decomposition of limestone during the manufacture of cement. The increasing amount of carbon dioxide in the air is mainly responsible for global warming.

Ø Global Warming and Greenhouse Effect: Some gases like carbon dioxide, methane, ozone, water vapours, CFCs have the capacity to trap some of the heat radiations that are released from the earth or from sun. These gases are known as greenhouse gases and the effect is called greenhouse effect. This leads to global warming.

Consequences of Global Warming: (i) It leads to melting of polar ice caps and flooding of low lying areas all over the earth. (ii) Global rise in temperature increases the incidence of infectious diseases like dengue, malaria, yellow fever,

sleeping sickness etc. Ø Acid Rain-When the pH of the rain water drops below 5.6, it is called acid rain. Oxides of nitrogen and sulphur

released as a result of combustion of fossil fuels dissolve in water to form nitric acid and sulphuric acid.

2SO2(g) + O2(g) + H2O(l) → 2H2SO4(aq)

4NO2(g) + O2(g) + H2O(l) → 4HNO3(aq)

Harmful Effects of Acid Rain: (i) It has harmful effects on trees and plants as it dissolves and washes away nutrients needed for their growth. (ii) It has very bad effect on aquatic ecosystem. (iii) Acid rain damages buildings and other structures made of stone or metal. Taj Mahal in India has been affected

by acid rain. (ii) Particulate Pollutants– Particulate pollutants are the minute solid particles or liquid droplets in air. Particulates in

the atmosphere may be viable or non-viable. Viable Particulates: They are minute living organisms that are dispersed in the atmosphere. e.g., bacteria, fungi,

moulds, algae etc. They cause plant diseases. Non-Viable Particulates: It can be classified according to their size and nature as follows: (i) Smoke: It is the mixture of solid and liquid particles formed during combustion of organic matter. Example:

cigarette smoke, smoke from burning of fossil fuel. (ii) Dust: It is composed of fine solid particles (over 2 gm in diameter). It is produced during crushing, grinding

and attribution of solid particles. (iii) Mist: These are produced due to the spray of liquids like herbicides and pesticides over the plants. They

travel through air and form mist. (iv) Fumes: They are generally released to the atmosphere by the metallurgical operations and also by several

chemical reactions. Harmful Effects of Particulate Pollutants: (i) Fine particles less than 5 microns penetrate into the lungs. Inhalation of such particles can lead to serious

lung diseases including lung cancer. (ii) Suspended particles of bigger size can hinder the sun rays from reaching the earth surface. This can lower

the temperature of earth and make the weather foggy. Ø Smog: This is the common form of air pollution which is combination of smoke and fog. Smog exists in two types: (i) Classical Smog: It occurs in cool humid climate. It contains smoke, fog and sulphur dioxide. It is also called

as reducing smog. (ii) Photochemical Smog: This type of smog results from the action of sunlight on unsaturated hydrocarbons

and nitrogen oxides released by the vehicles and industries. It has high concentration of oxidising agents and is therefore, called as oxidising smog.

Formation of Photochemical Smog-

NO2(g) hv → NO(g) + O(g)

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 101 Oxygen atoms are very reactive and combine with the O2 in air to produce ozone.

O(g) + O2 (g) � ⇀�↽ �� O3 (g)

NO (g) + O3 (g) → NO2 (g) + O2 (g)

3CH4 + 2O3 → 3CH2 = O + 3H2O

Harmful effects of Photochemical Smog: (i) They damage metals, stones, building materials etc. (ii) They produce irritation in the eyes and respiratory system Control of Photochemical Smog: l To control this type of pollution, the engines of the automobiles are fitted with catalytic converters to check

the release of both oxides of nitrogen and hydrocarbons in the atmosphere. l Some plants like Vitis, Pinus, Juniparus, Quercus, Pyrus can metabolise nitrogen oxide and therefore, their

plantation can be done. Ø Stratospheric Pollution l Formation of Ozone: Ozone in the stratosphere is produced by UV radiations. When UV – radiations act on

dioxygen (O2) molecules, ozone is produced. Ozone is thermodynamically unstable and decomposes to molecular oxygen. Thus, there exists equilibrium between production and decomposition of ozone molecules.

l Depletion of Ozone layer: Ozone blanket in the upper atmosphere prevents the harmful UV radiations from reaching earth. But in recent years, there have been reports of depletion of this layer due to presence of certain chemicals in the stratosphere. Chlorofluorocarbons (CFCs), nitrogen oxides, CCl4 etc. are the chemicals responsible for depletion of ozone layer.

Cl CF C l CClF2 2 2 → +

Cl ClO

O O+ → +3 2

ClO O Cl

+ O+ → 2

Chlorofluorocarbons dissociate in the presence of light give chlorine free radicals which catalyse the conversion of ozone into oxygen.

lEffects of the Depletion of Ozone Layer: (i) This leads to many diseases like skin cancer, sunburn, ageing of skin, cataract etc. (ii) UV radiations can kill many phytoplanktons, damage the fish productivity. (iii) It can decrease moisture content of the soil by increasing the evaporation of surface water. (iv) UV radiations can damage paints and fibres, causing them to fade faster.

Know the TermsØ Environmental Studies: It deals with the sum of all social, economical, biological, physical and chemical

interrelations with our surroundings. Ø Troposphere : The lowest region of atmosphere in which the human beings along with other organisms live is

called troposphere. Ø Greenhouse Effect : The progressive warming up of the earth’s surface due to blanketing effect of man-made

carbon dioxide is called greenhouse effect. Ø Acid Rain: When the pH of the rain water drops below 5.6, it is called acid rain.

Know the Facts Ø The effects of particulate pollutants are largely dependant on the particle size. Air borne particles are dangerous

for human health. Ø Two kinds of pollutants occur- (i) Primary Pollutants: Pollutants which are emitted directly from the sources. (ii) Secondary Pollutants: Pollutants which are formed in the atmosphere by chemical interaction among

primary pollutants.

HC3CO

O

ONO2

Peroxyacetyl nitrate (PAN)

CH2= HC CH = O

Acrolein


TOPIC-2Water and Soil Pollution, Strategies to Control Environmental Pollution

Quick Review Ø Water Pollution: Presence of undesirable materials in water which is harmful for the human beings and plants is

known as water pollution. Causes of Water Pollution: Pathogens: The most serious water pollutants are the disease causing agents called pathogens include bacteria

and other organisms that enter water from domestic sewage and animal excreta. Organic wastes: The other major water pollutant is organic matter such as leaves, grass, trash etc. Excessive

phytoplankton’s growth within water is also a cause of water pollution. Ø BOD (Biochemical Oxygen Demand): It is defined as the amount of oxygen required by bacteria for the

breakdown of the organic matter present in a certain volume of a sample of water. Ø The amount of BOD in water is a measure of the amount of organic material in the water. Clean water has BOD

value of less than 5 ppm. Highly polluted water could have a BOD value of 17 ppm or more. Ø Chemical Pollutants: Water soluble heavy metals such as cadmium, mercury, nickel etc. constitute an important

class of pollutants. All these metals are dangerous to humans because our body can’t excrete them. The organic chemicals are another group of substances that are found in polluted water. Petroleum products

pollute many sources of water. Ø International Standards for Drinking Water: Fluoride: Concentration of fluoride above 2 ppm causes brown mottling of teeth. Excess of fluoride is harmful to

bones also. Lead: Upper limit concentration of lead in drinking water is about 50 ppm. Lead can damage kidney, liver,

reproductive system etc. Sulphate: Excessive sulphate in drinking water causes laxative effect, otherwise at moderate levels it is harmless. Nitrate: The maximum limit of nitrate in water is 50 ppm. Excess nitrate can cause methemoglobinemia (“blue

baby syndrome”). Ø Soil Pollution: Soil pollution or soil contamination is caused by the presence of xenobiotic (human made)

chemicals or other alterations in the natural soil and environment. Ø Causes of Soil Pollution: l Waste Dumping: Industrial solid wastes and sludge are the major sources of soil pollution. l Mining: The vast variety of toxic chemicals released by mining activities can harm animals and aquatic life as

well as their habitat. l Pesticides: Many of the chemicals used in pesticides are persistent soil contaminants, which adversely affect

soil conservation. Ø Industrial Wastes: These are also sorted out as non-degradable wastes which are generated by thermal power

plants which produce fly ash, integrated iron and steel melting slag. The disposal of non-degradable industrial solid wastes, if not done by proper method, may cause serious threat to the environment.

Now-a-days, fly ash and slag from the steel industry are utilised by the cement industry. Ø Strategies to control environmental pollution: (i) The improper disposal of wastes, reuse is one of the major causes of environmental degradation. Waste

management i.e., reduction of the waste, reuse and proper disposal wherever possible recycling of materials and energy must be done.

(ii) All domestic wastes should be properly collected and disposed. Ø Green Chemistry: Green chemistry is about utilising the knowledge and principles of chemistry that would

control the increasing environmental pollution. Ø Green chemistry in day-to-day life: (i) Dry-Cleaning of clothes and laundary: Replacement of halogenated solvent like (CCl4 ) by liquid CO2 which

is less harmful to groundwater. Hydrogen peroxide (H2O2 ) is used for the purpose of bleaching clothes. (ii) Bleaching of Paper: In place of chlorine, H2O2 is used for the bleaching of paper. (iii) Synthesis of Chemicals: Ethanal (CH3CHO) is prepared by step oxidation of ethene. Such as, CH2 = CH2 + O2 → CH3CHO in the presence of catalyst Pd, Cu [in water]

Oswaal CBSE Chapterwise Quick Review, CHEMISTRY, Class-XI [ 103Know the Terms

Ø Pathogens: The most serious water pollutants are the disease causing agents called pathogens. Ø Dissolved Oxygen: The amount of oxygen that water can hold in the solution is called dissolved oxygen. Ø Green Fuel: Fuel obtained from plastic waste has high octane rating. It does not contain lead and is called “green

fuel”.

Know the Facts Ø Eutrophication: The process in which nutrient enriched water bodies support a dense plant population, which

kills animal life by depriving it of oxygen and results in subsequent loss of biodiversity is known as eutrophication. Ø COD (Chemical Oxygen Demand): It is calculated as the amount of oxygen required to oxidise the polluting

substances. It is measured by treating the given sample of water with an oxidising agent, generally K2Cr2O7 in the presence of dil. H2SO4 .

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Documents

General Introduction and Nature of Matter Quick Review · Chapter - 1 : Some Basic Concepts of Chemistry TOPIC-1 General Introduction and Nature of Matter Quick Review Ø Chemistry